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DISCLAIMER! ▪ I am not vouching for the accuracy of any of the information presented in the following slides. ▪ I am sharing the work of your peers so that you may have additional explanations of the material to benefit you. ▪ All the slides pertaining to a single topic are grouped together. SCIENTIFIC METHOD You have completed kindergarten at [North Hills Preparatory] and are awarded this diploma in recognition of your accomplishments. How to memorize? Angry Hippos Eat Dry Cereal What does it mean? Angry – Ask a testable question Hippos – Form a reasonable hypothesis Eat - Experiment Dry – Data Analysis Cereal – Form a conclusion SCIENTIFIC METHOD By: Keiji Chan and Nael Alami The Basics The Scientific Method is the logical process to solving problems by: 1. Observing 2. Collecting data 3. Creating hypotheses 4. Testing hypotheses 5. Formulating theories backed up by data • The scientific method can also be a continuous cycle Observing & Collecting Data • Also known as research Quantitative Data • Numerical • Measured • Example: There are 7 birds outside. Qualitative Data • Describes things • Not in number form • Example: The birds are blue. • Data does not have to be only numbers, but descriptive ideas too. Formulating & Testing Hypotheses • A statement that provokes further investigation and can be proven right or wrong • Can be written in an “if…then” format, a form of cause and effect statement • Test one variable, preferably – Ensures more definite results • Must be tested multiple times for accuracy • Based off of some sort of observation or past evidence • Even if proven incorrect, a hypothesis should be revised, not rejected Controls & Variables Controls • Factors not tampered with during an experiment • Examples: – Moisture, light exposure, most environmental conditions Independent Variable • The factor being changed for the purpose of the experiment Dependent Variable • Factor being measured • Depends on independent variable Models & Theories Models • Explanation of the process and connections of an occurrence – Can come in many forms (ex: diagram) – Small scale version of something that proves a theory Theories • Basic overview of the facts supporting a conclusion • Can be represented in models • Can foresee the results of an experiment SCIENTIFIC METHOD JOSUE AND FARLEY HYPOTHESIS • IS A EDUCATED GUESS • SERVES AS BASIS FOR MAKING PREDICATIONS • ALSO CARRYING OUT FURTHER EXPERIMENTS(LEADING TO THEM). CONTROLS • EXPERIMENT DESIGNED TO MINIMIZE THE EFFECTS OF VARIABLES OTHER THAN THE INDEPENDENT VARIABLE. • IN OTHER WORDS, DOING AN EXPERIMENT THE SAME BUT WITHOUT VARIABLES THAT YOU'RE TESTING FOR, SO YOU CAN COMPARE RESULTS. • E.G: SEE HOW WELL CAFFEINE WILL AFFECT YOUR BODY. THE ONE WITHOUT CAFFEINE IN THE THEIR BODY IS THE CONTROL. • WANTS TO REMAIN THE SAME VARIABLES • INDEPENDENT VARIABLE -IT'S MANIPULATED OR IT IS CHANGED BY THE SCIENTIST. • DEPENDENT VARIABLE -IT IS OBSERVED OR MEASURED IN THE EXPERIMENT. • A VARIABLE IS ANY ITEM OR FACTOR OR CONDITION THAT CAN BE CONTROLLED OR CHANGED IN THE EXPIREMENT. THEORIES • THEORY- A GROUP OF PROPORTIONS CREATED TO EXPLAIN A GROUP OF FACTS OR PHENOMENA. • THEORIES IN THE EXPERIMENT ARE PREDICTIONS OF WHAT IS THOUGHT TO BE THE OUTCOME. • A THEORY IS CONSIDERED TO BE CORRECT TO BE TRUE/ USEFUL IF IT CAN PREDICT THE RESULTS OF NEW EXPERIMENTS. The Scientific Method Thomas Arnett Steps • Ask a Testable Question • Make a good hypothesis • Do the experiment • Analyze Data • Make a conclusion • Angry Hippos Eat Dried Cereal Question • Needs to be testable • Not a yes or no question • Should be what your experiment is based on Formulating a hypothesis • The hypothesis is a short statement of what you will be doing and what you think will happen • Should be testable as well • Does not need to be an if…then statement Experimenting • The experiment needs to be easily replicable • You will need: • Procedures • Materials • In addition to those, you need to observe data • Data can be qualitative or quantitative Data Analysis • This is where you review your data, do your math, and eventually begin to conclude • Should explain what your observations and data mean and why they are significant • Needs to analyze everything recorded Conclusion • Conclude if your hypothesis was correct or not • Explain why this information is relevant • Should use your data to back up your concluding statements MEASUREMENTS MEASUREMENTS, SIGNIFICANT FIGURES, SCIENTIFIC NOTATION, PRECISION AND ACCURACY By Manasi Taduri & Sampreeti Bingi METRIC UNIT MEASUREMENTS • The IS (International System of units) is used all around the world. Except America. • It was created in 1799 and was used temporarily to replace the previously existing system and then replaced it for good since it as easier to understand. Base Measurement Symbol Customary Equivalent Meter Length m Feet Gram Mass g Ounces Liter Volume L Fluid ounces Second Time s second PREFIXES LESS THAN THE BASE • Deci – 10^-1 x base; 1/10 base units • Centi – 10^-2 x base; 1/100 base units • Milli - 10^3 x base; 1/1,000 base units • Micro - 10^6 x base; 1/1,000,000 base units • Nano - 10^-9 x base; 1/1,000,000,000 base units • Pico - 10^-12 x base; 1/1,000,000,000,000 base units PREFIXES CONT. GREATER THAN THE BASE • Deca – 10^1 x base; 10 base units • Hecto – 10^2 x base; 100 base units • Kilo – 10^3 x base; 1,000 base units • Mega - 10^6 x base; 1,000,000 • Giga – 10^9 x base; 1,000,000,000 • Tera – 10^12 x base; 1,000,000,000,000 SIGNIFICANT FIGURES • All number 1 through 9 are significant • Zeros are NEVER significant when at the beginning of a number; they are only place holders. • Example: 0.000748 only has three sig figs. • Zeros in the middle of the number are ALWAYS significant. • Example: 9087 has four sig figs • Zeros are only significant at the end of a number if there is a decimal point. • Examples: 7850 has three sig figs. 8490.0 has five sig figs. 0.09840 has four sig figs. SCIENTIFIC NOTATIONS • Scientific notation is a way to easily convert very large numbers or very small numbers to something more manageable. • For numbers less than zero, you need to move the decimal to the right until you hit the SECOND non-zero number. • For example, if you have .000000312, you would move the decimal 7 numbers to the RIGHT so you end up with 3.12 and then you add the power of 10 which is the number of spaces you moved (negative exponent if you moved right) so you would get 3.21 x 10^-7. • For numbers larger than zero, you do the same thing with one difference. You move the decimal to the left until your decimal is behind the first non-zero digit in the number. • Example: If you have 456000000. then you move the decimal 8 places to the LEFT so you get 4.56 and then you add the power of ten and it would be 10^8 so your final answer would be 4.56 x 10^8. :) ACCURACY AND PRECISION • Accuracy: the degree to which the result of a measurement, calculation, or specification conforms to the correct value of the correct value or a standard • This means that the values are close to the intended target but not close to each other. • Precision: refinement in a measurement, calculation, or specification, especially as represented by the number of digits given MEASUREMENT By: Alea Badayos & Serena Myoung BASE UNITS Meter units Quantity Abbreviation Customary units Meter Length m foot Gram Mass g Ounces Liter Volume L Fluid ounces Second Time s seconds PREFIXES Prefix Abbv. Exponential form Tera T 10^12 x base giga G 10^9 x base mega M 10^6 x base kilo K 10^3 x base Hecto h 10^2 x base Deca D 10^1 x base BASE (m/L/g/s) Deci d 10^-1 x base Senti c 10^-2 x base Milli m 10^-3 x base Micro µ 10^-6 x base Nano n 10^-9 x base pico p 10^-12 x base UNIT CONVERSIONS Prefix Base Units 1. Replace prefix with corresponding integer power of ten. 2. Simplify value (move decimal the correct number of places in the correct direction). EX: 354.84 milliliter 354.84 x 10 -3 L “10 -3 ” means the decimal moves 3 places to the LEFT 354.84 milliliter s = 0.35484 liters Negative power of ten = decimal is moved to the left Positive power of ten = decimal is moved to the right UNIT CONVERSION Base Units Prefix 1. Divide the measurements by its corresponding integer of power of ten. 2. Add prefix name to the unit. If multiplication is preferred, instead of division … 1. Apply -1 exponent to power of 10 to move denominator to numerator, so it is easier to solve.. 2. Multiply measurement by new power of ten. EX: meters kilometers (kilo= 10 3 ) 1609 meters/10 3 1609 x 10 -3 1 .609 kilometers Negative exponent = move the decimal to the left. UNIT CONVERSION Prefix Prefix 1. Multiply the number by its own integer power of 10. 2. Divide the second prefix’s integer power of 10 last. EX: 818 mL µL; m=10 -3 , µ=10 -6 818 x (10 -3 /10 -6 ) x 10 6 10 3 818 x 10 3 = 818,000µL Add the exponents ACCURACY & PRECISION Accuracy: the closeness of the measurements to the accepted value (your “bulls eye”) of the quantity measured. Precision: closeness of the set of measurements of the same quantity; the numbers are close to eachother. SIGNIFICANT FIGURES (SIG FIGS) Digits that carry meaning contributing to the exactness of the measurements. Consist of all digits known with certainty plus one final digit that is estimated. Rules of sig for 0: Beginning of the number: NEVER Middle of the number: ALWAYS End of the number: SOMETIMES - When the zero has non-zero digits before it (e.g.0.650): YES - When there’s a decimal present after the zero (e.g. 650.0): YES - When there’s no decimal present after the zero (e.g. 650): NO SCIENTIFIC NOTATION Large numbers written in an easy format and is rounded to the hundredth place: #.## x 10 n When the exponent is positive, the decimal place is moved to the right, making the number greater than one. When the exponent is negative, the decimal place is moved the left, making the number less than one. EX: 0.0000076392 7.64 x 10 -6 4315.09 4.32 x 10 3 0.00054925 x 10 4 5.4925 x 10 4 Measurements By Waris Sandhu and John Henry Nwachukwu International System of Units • • • System of measurements based on prefixes changing the amount of the base Bases include grams, liters, meters Prefixes change the amount of base in integers of 10 Base Units • Grams measure mass (g) • Meters measure distance (m) • Seconds measure time (s) • Liters measure volume (L) Prefixes • Pico= 10-12 • Nano=10-9 • Micro=10-6 • Milli=10-3 • Centi=10-2 • Deci=10-1 To get the value of a base with a prefix multiply the base by the values. Prefixes Contd. • Tera= 1012 • Giga= 109 • Mega= 106 • Kilo= 103 • Hecto= 102 • Deca= 101 Prefix Abbreviations • Pico (p) • Nano (n) • Micro (µ) • Mega (M) • Giga (G) Prefix Abbreviations Cont. • Kilo (k) • Hecto (h) • Deca (D) • Deci (d) • Centi (c) • Milli (m) Measuring with Metric Rulers • Measurements on a ruler should be estimated one digit larger that the certain measurement. (if the measurement is .4 centimeters, it should be written as .40 cm) Graduated Cylinder • Meniscus- A curve at the top of a liquid caused by surface tension • Measuring should be done from the bottom of the meniscus Accuracy and Precision ACCURACY • How close you are to the true value. Precision • How close your measurements are to each other. Sig Figs • Significant Figures- Digits that contribute exactness towards the measurement. • Measurement where all digits are certain except for the last which is rounded or estimated. Counting Sig Figs • All non-zero numbers are sig figs • Zeroes at the beginning of a number are never sig figs • Zeroes in the middle of the number are always sig figs • Zeroes at the end of a number are sometimes sig figs • • When there is a decimal present it is a sig fig When there is no decimal it is not a sig fig Counting Sig Figs contd. • Whole quantities, Constants, and Conversion Factors have infinite sig figs. • To find sig figs after adding and subtracting, the larger of the smallest place value is used to determine what the place value of the final sig fig will be in the result. • To find sig figs after multiplying or dividing, the total sig figs in the product is equal to the least number of total sig figs of the numbers that are being divided or multiplied. Scientific Notation • Used for writing really small or large numbers into forms that are easier to read • Positive exponent, number is greater than one • Negative exponent, number is less than one • Number should always be to the hundredths place Unit Conversions • To convert to base unit, replace the prefix with its corresponding integer power of ten. Then, simplify the value by moving the decimal the appropriate number of times in the correct direction. • 400.00 millimeters > .4 meters Converting from the Base Unit • To convert from base to prefix-unit, divide by the corresponding integer power of ten and add the correct prefix name to the unit • 1700 meters/ 103 > 1.7 kilometers ATOMS Atoms AT O M - T H E BA S I C U N I T O F A C H E M I C A L E L E M E N T Early Atomic Theory 1766-1844 Aristotle Democritus Believed in the four elements of Believed matter could only be air, water and fire Believed matter could be divided infinitely without changing it’s properties divided until you got to the smallest particle (atom) Atoms join with out Plato Believed the world was created from four elements (air, earth, fire, water) Modern Atomic Theory 1913 J.J. Thompson Discovered the Electron Plum Pudding Model (1904) Identified the negativity charged electron in the cathode ray tube in 1897 Plum Pudding Model According to this model, the atom is a sphere of positive charge, and negatively charged electrons are embedded in it to balance the total positive charge. Cathode Ray Tube A cathode ray tube (CRT) is a specialized vacuum tube in which images are produced when an electron beam strikes a phosphorescent surface. Ernest Rutherford Proposed the nuclear atom as the result of the gold-foil experiment in 1911 Concluded that the atom is mostly empty space Gold-Foil Experiment Robert Millikan A famous physicist The oil drop experiment was the most famous scientific work of his career Goal: to quantify the charge of an electron Law of conservation of mass “Mass is neither destroyed nor created during ordinary chemical reactions or physical changes.” Law of definite proportions “The fact that a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound” Law of multiple proportions “When two elements combine to form more than one compound, the mass of one element, which combines with a fixed mass of the other element, will always be ratios of whole numbers.” ATOMIC THEORY BY Z.C P.S ATOMIC THEORY JOKES ▸ funny jokes Thats Punny. -Parker S.S. ATOMIC THEORY ▸ Definition: any theory in which matter is regarded as consisting of atoms. ‣ T current theoretical model of the atom involves a dense nucleus surrounded by a probabilistic "cloud" of electrons. In chemistry and physics, atomic theory is a scientific theory of the nature of matter, which states that matter is composed of discrete units called atoms. ATOMIC THEORY HISTORY LESSON ‣ over the centuries there have been multiple versions of Atomic Theory. Early Atomic Theory, Early Modern Atomic Theory, Dalton’s Atomic Theory, and Modern Atomic Theory. ‣ The ancient Greek philosopher Aristotle he postulated that all mater was made of earth, water, air and fire ‣ in 1789, Antoine Lavoisier formulated the total mass of reactants and products in a chemical reaction remains constant ATOMIC THEORY FOUNDERS OF THE ATOMIC THEORY ▸ Democritus who predicted that everything is composed of atoms and empty void. ▸ Plato believed that the world was created from four elements ▸ Aristotle argued that the elements were not made of atoms ▸ Antoine Lavoisier formulated that total mass of reactants and products in a chemical reaction remains constant ▸ John Dalton created a Atomic Theory that helped support Democritus’s prediction ▸ Nelis Bohr’s model of the atom is important because it introduced the concept of the quantum in explaining atomic properties. ATOMIC THEORY LAWS OF THE ATOMIC THEORY ▸ law of conservation of Mass ,1789 Mass is neither destroyed nor created during ordinary chemical reactions or physical changes. ▸ law of Definite Proportions, 1799 A chemical compound contains the same elements in exactly the same proportions by mass regardless of the size. ▸ law of multiple proportions, 1808 if two or more different compounds of the same two elements, then the ratio of the masses of the second element combined with the first element is always a ration of small whole numbers. THE DIFFERENT ATOMIC THEORIES ▸ Early Atomic Theory: Proposed that everything is composed of atoms and empty voids. ▸ Early Modern Atomic Theory: Is the total mass of reactants and products in a chemical reaction remains constant. ▸ Dalton’s Atomic Theory: Dalton’s is the first complete attempt to describe all matter in terms of atoms and their properties. ▸ Modern Atomic Theory: is a theory that all matter is composed of tiny particles called atoms. FUN FACTS: ▸ “Atoms comes from Greek word “Atomos,”meaning “indivisible.” ▸ In 1897, J.J. Thompson created the Cathrode Ray Tube that discovered the Electron ▸ In 1909, Ernest Rutherford performs the gold foil experiment to discover the nucleus. ▸ In 1932, James Chadwick discovered the Neutron BIBLIOGRAPHY: ▸ https://www.reference.com ▸ www.iun.edu/~cpanhd/.../modern-atomic-theory/modern-atomic-theory. ▸ https://www.khanacademy.org/.../atomic-structure.../daltons-atomic-theory-version- Atoms By Hrishikesh and Tarun Early Atom Theory Democritus ➔ Early 5th century BCE ➔ Proposed everything is made of atoms and empty void ➔ Atom comes from the word, Atomus meaning indivisible Plato ➔ Late 5th century BCE ➔ Believed the world was made from 4 elements: Water, Earth, Fire, and Air Aristotle ➔ Approx: 330BCE ➔ Argued against atoms, saying that the elements were not made of atoms ➔ Also believed change happened through transformation instead of rearrangement Modern Atomic Theory Further discoveries led to 2 postulates of Dalton’s atomic theory being discredited. ➔ Atoms can be split into 3 subatomic particle- the electron (1e), neutron (neutral) and the proton (+1e). ➔ Atoms of same element can have different masses as proved by the discovery of Isotopes Law of Conservation of Mass In 1789 ➔ Mass cannot be destroyed nor created in chemical reactions ➔ Every chemical reaction has an equal mass for reactants and products. Formulated by Antoine Lavoisier Law of Definite Proportions In 1799 ➔ A chemical compound contains the same elements in the exact same proportion by mass, regardless of the source or size of the sample Proved by Joseph Louis Proust Law of Multiple Proportions In 1808 ➔ States that when two elements combine to form more than one compound, the mass of one element, which combines with a fixed mass of the other element, will always be ratios of whole numbers. ➔ Applies only to compounds composed of the same elements Proposed by John Dalton Subatomic Particles Electron: ➔ Extremely small negatively charged particle that orbits the nucleus ➔ AMU ≈ 0 ➔ Discovered by J.J. Thompson in the cathode ray experiment Charge: -1e Proton ➔ Positively charged particle ➔ AMU = 1 ➔ DIscovered by Ernest B. Rutherford Charge: +1e Neutron ➔ Particle roughly the same size as a proton with no charge ➔ AMU =1 Charge: 0 (neutral) Isotopes ➔ Concept first proposed by F. Soddy. ➔ Atoms of the same element that contain different masses due to different number of neutrons ➔ Unstable isotopes are radioactive. ➔ Isotopes can be represented in 2 ways: helium-4 or ⁴He ( With 4 being the mass number of that respective isotope). Average Atomic Mass ➔ Calculated by summing the mass of the element’s isotopes, each multiplied by its natural abundance on Earth. ➔ Average atomic mass = f1M1 + f2M2 + ... + fnMn where f is the fraction representing the natural abundance of the isotope and M is the mass number (weight) of the isotope. Example: Average atomic mass of chlorine 1st isotope of chlorine= 35 amu, 75.77%, 2nd isotope= 37 amu, 24.23% (0.7577 * 35 amu) + (0.2423 * 37 amu) = 35.483 amu Bibliography: http://www.ducksters.com/science/chemistry/isotopes_hydrogen.gif http://images.tutorvista.com/cms/images/44/Atom.png https://upload.wikimedia.org/wikipedia/commons/thumb/5/56/Law_of_Definite_Proportions.svg/2000pxLaw_of_Definite_Proportions.svg.png http://study.com/cimages/multimages/16/nitrogenoxygencompound2.png Subatomic Particles, Isotopes, Avg. Atomic Mass By: Tejas Saboo, Nick Patel, and Mariana Zollinger Subatomic Particles Protons Located: in the nucleus Charge: Positive (+) Mass: 1 amu Discoverer: Ernest Rutherford Neutrons Located: in the nucleus Charge: No Charge (0) Mass: 1 amu Discoverer: James Chadwick Subatomic Particles Electron Located: in the electron cloud Charge: Negative (-) Mass: (basically) 0 amu Discoverer: J. J. Thompson Isotopes ● Atoms of the same element have different masses,resulting in them having different mass numbers -isotopes can be written as the element name followed by the mass number The mass of the atom changes if the number of neutrons change Some isotopes are more stable than others and are commonly found in nature Average Atomic Mass Atomic Mass Unit (AMU) is exactly 1/12 the mass of a Carbon-12 Atom Average Atomic Mass= Average of the atomic masses of the different isotopes of a single element (based on their percent of natural abundance) Natural Abundance= Abundance of isotopes of an element naturally found on a planet Calculating the average atomic mass of an element is just like calculating a weighted average! HOW TO CALCULATE YOUR CHEM GRADE CATEGORY GRADE X WEIGHT = CONTRIBUTION TO AVG. Minor Work 100 X 20% = 20 Minor Lab 70 X 30% = 21 = 45 = 86 In calculating the Average Atomic Mass of an element: Test 90 X 50% Category= Isotope Grade= Atomic Mass (AMU) Weight= Natural Abundance TOTAL Subatomic Particles, Isotopes, and Average Atomic Mass By: Kumar Kudumula and Zain Abbas The Electron Discovered by Joseph John Thomson a English physicist through the use of the cathode ray tube experiment . Conclusions: Cathode Rays are negatively charged Cathode ray particles are 2000 times lighter than the hydrogen Ion Atoms must have positive and negative charges Atomic Model after the Discovery of the Electron The Plum Pudding Model Visualizes Thomson's last conclusion that atoms are positive spheres with electrons scattered throughout Electrons Cont. Robert A. Millikan preformed the oil drop experiment in 1909. Drop of oil moves up or down based on the voltage of the metal plate Conclusion: The charge of an electron is (-1e) or -1 times the elementary charge which is 1.602 x 10-19 Coulombs Discovery of the Nucleus Ernest Rutherford preformed the gold foil experiment in 1909. Fired a beam of positively charged particles at a sheet of gold foil Most particles went through the foil but 1 in 20,000 were deflected back Rutherford’s Conclusions: The positive charge of an atom is concentrated in one place Atoms have a dense, positively charged center Rutherford's Conclusions Cont. Electrons are concentrated in a cloud around the nucleus Mass of an element comes from the nucleus Different elements have differently sized nuclei Electrons are made of positively charged particles, which are known as protons Protons have a mass of 1 amu and have a charge of +1 time the elementary charge Rutherford’s Model of the Atom A central Nucleus made up of positively charged protons Electrons existing in a cloud around the nucleus The Neutron James Chadwick shoots Beryllium with alpha particles from Polonium. Radioactive particles with no charge are emitted from the nucleus. Chadwick’s Conclusions: The uncharged particles are now known as Neutrons Neutrons have a mass of 1 amu Neutrons stabilize the nucleus Isotopes Atoms of the same element that have different masses Can be written as the element name followed by the mass number Ex: Titanium with a mass of 58 Titanium-58 Mass of an atom changes if the total amount of neutrons varies between each atom Average Atomic Mass Definition: The weighted average of the atomic masses of the naturally occurring isotopes of an element. Atomic Mass: An atomic mass unit is 1/12 of the mass of a carbon-12 atom The atomic mass of any nuclide is determined by comparing it with the mass of a Carbon-12 atom. Finding Average Atomic Mass In order to find the average atomic mass of an element you need both the atomic masses of its isotopes and their natural abundance. Finding average atomic mass of Silver: Step 1: Find the masses of each isotope Step 2:Find the abundance of each isotope Step 3: Turn the percentages of each isotopes into decimals Step 4: Multiply the atomic mass of each isotopes with their respective isotope Finding Average Atomic Mass of Ag Natural Abundance Contribution to the Average Isotope Atomic Mass Ag-107 106.950 Amu X 0.5186 = 54.441 amu Ag-109 108.904 Amu X 0.4814 = 52.4267 amu Finally add both numbers together to get your average atomic mass which is 107.868 atomic mass units. SUBATOMIC PARTICLES, ISOTOPES, AND AVERAGE ATOMIC MASS SHEREBANOO & NIBITIKA SUBATOMIC PARTICLES • PARTICLES THAT ARE SMALLER THAN AN ATOM. • ALSO KNOWN AS ELEMENTARY PARTICLE. • INCLUDES: PROTONS, NEUTRONS, AND ELECTRONS. ELECTRONS • SUBATOMIC PARTICLES THAT HAVE A NEGATIVE CHARGE. • REPRESENTED BY THE - SIGN • FOUND IN THE ELECTRON CLOUD, OUTSIDE THE NUCLEUS. • DISCOVERED BY JJ THOMSON THROUGH THE CATHODE RAY TUBE EXPERIMENT. • THOMSON CONCLUDED: 1. PARTICLES IN THE CATHODE RAY WERE 2000 TIMES SMALLER THAN THE LIGHTEST PARTICLE AT THE TIME, WHICH WAS THE HYDROGEN ION. 2. ATOMS ARE POSITIVE SPHERES WITH ELECTRONS SCATTERED THROUGHOUT. PLUM PUDDING MODEL PROTONS • SUBATOMIC PARTICLES THAT HAVE A POSITIVE CHARGE. • REPRESENTED BY THE + SIGN. • FOUND IN THE NUCLEUS OF ATOM. • DISCOVERED BY ERNEST RUTHERFORD THROUGH THE GOLD FOIL EXPERIMENT. • RUTHERFORD’S DISCOVERIES: 1. POSITIVE CHARGE IS FOCUSED ON ONE PART OF THE ATOM. 2. THAT ONE PART IS CALLED THE NUCLEUS. 3. ELECTRONS ARE IN A CLOUD AROUND THE NUCLEUS. RUTHERFORD’S MODEL OF THE ATOM NEUTRONS • SUBATOMIC PARTICLES THAT HAVE A NEUTRAL CHARGE. • FOUND IN THE NUCLEUS. • FOUNDED BY JAMES CHADWICK IN 1932. • CHADWICK BOMBARDED BERYLLIUM WITH ALPHA PARTICLES. • HE DISCOVERED: 1. NEUTRONS STABILIZE THE NUCLEUS AND KEEP PROTONS FROM REPELLING EACH OTHER. 2. NEUTRONS HAVE A MASS OF ABOUT 1 AMU. (AN ATOMIC MASS UNIT) ISOTOPES • ATOMS OF THE SAME ELEMENT THAT HAVE DIFFERENT MASSES, THEREFORE DIFFERENT MASS NUMBERS. • SAME NUMBER OF PROTONS BUT DIFFERENT NUMBER OF NEUTRONS. • IDENTICAL CHEMICAL PROPERTIES (REACTIVITY, FLAMMABILITY, ETC.) BUT DIFFERENT NUCLEAR PROPERTIES (RADIOACTIVITY, ETC.) AVERAGE ATOMIC MASS • THE WEIGHTED AVERAGE OF THE ATOMIC MASSES OF THE NATURALLY OCCURRING ISOTOPES OF AN ELEMENT. • NATURAL ABUNDANCE: THE PERCENTAGE OF THE ISOTOPE OF ALL NUCLIDES FOR THAT GIVEN ELEMENT. NUCLIDE IS A GENERAL TERM FOR ANY ISOTOPE OF ANY ELEMENT. FOUND ON THE PERIODIC TABLE: THE ATOMIC MASS Atoms WITH YA GIRLS LEI AND GEN What is an atom? Atoms are the #basic units of matter. They make up everything. >.> those liars… They’re made up of THREE particles: Electrons. Protons, and Neutrons Pre- Info Protons’ Mass≈Neutron’s Mass One of either>1800 electrons Atoms… Have equal protons and electrons Adding a proton to an atom makes a new element, whereas adding a neutron makes an isotope PROTONS Positively Charged Found in Nucleus Ernest Rutherford Number of Protons=Atomic # determines what element it is and chemical behavior EX. H:1, C:6, O:8 Electrons Negatively charged Electrically attracted to protons Electron configuration: orbital description of the locations of the electrons in an unlit/unexcited atom –helps chemists predict boiling point, stability, and conductivity Ex. Be:1s^2,2s^2, also written as Be:{HE}2s^2 SPDF Electron Configuration Electron configuration: orbital description of the locations of the electrons in an unlit/unexcited atom –helps chemists predict boiling point, stability, and conductivity Ex. Be:1s^2,2s^2, also written as Be:{HE}2s^2 SPDF: SIMPDF Interactive.. Activity… Cu He Ar Mg Se Neutrons • • • • Uncharged particles Slightly larger than a proton James Chadwick in 1932 Aren’t named or look like Jimmy ISOTOPES any of two or more forms of a chemical element, having the same number of protons in the nucleus, or the same atomic number, but having different numbers of neutrons in the nucleus, or different atomic weights. There are 275 isotopes of the 81 stable elements, in addition to over 800 radioactive isotopes, and every element has known isotopic forms. Isotopes of a single element possess almost identical properties. # of neutrons determines the isotope of that element Average Atomic Mass To find aam… Make note of the isotopes and their percentages The final calculation is a weighted average meaning it roughly the mass of the isotope times the relative abundance Add the results together AAM Activity Bibliography http://www.livescience.com/37206-atom-definition.html http://www.dictionary.com/browse/isotope ATOMIC STRUCTURES ATOMIC STRUCTURES SECTION: 4-1, 4-2, 5-3, 6-2 By: Macy and Kayleen Lights and Waves • Electromagnetic Radiation is a form of energy that exhibits wavelike behavior as it travels through space • All forms of electromagnetic radiation travel at a constant speed: (c= 3.00×108 m/s) • Wavelength (lambada λ) is the distance between corresponding points on adjacent waves • Frequency (v) is the number of waves that pass a given point in a specific time usually one second • Planck’s constant (h= 6.62607004 × 10-34 m2 kg / s) is the relationsip between quantuma nd frequency • C= λv E= hc/ λ E=hv Electromagnetic Spectrum Visible Light: 400 to 700 • Used to describe location of electrons in electron cloud Quantum Numbers • Four quantum numbers= complete set • Every electron has unqique set of numbers • Quantum numbers: n, ℓ, mℓ, ms • n= any whole number but 0 • ℓ= any number lower than n • mℓ = every number in between the negative value of ℓ and ℓ plus the negative value of ℓ and ℓ • ms = 1/2 or -1/2 Quantum Numbers and Orbitals • Orbitals: three dimensional region around the nucleus that indicates the probable location of an electron • Each l value corresponds with a different orbital shape • Each orbital shape can hold different amounts of electrons • S= sphere (1) • P= dumbbell (3) • D= Clover (5) L Value 0 1 2 3 4 5 6 Orbital S Max # of e 2 P 6 D 10 F 14 G 18 H 22 I 26 Lewis Dot Diagrams • Diagrams that show the bonding between atoms of a molecule and the lone pairs of electrons that may exist in the molecule • Created with the valence electrons of an element Electron Configuration • Is the arrangement of electrons in an atom • An electron always occupies the lowest energy orbital that can receive it (Aufbu principle) • No 2 electrons in same atom can have same set (Pauli exclusion principle) • Orbitals of equal energy are occupied by one electron before occupied by a second (Hund’s rule) Valence Electrons • Electrons lost, gained, or shared in chemical compound formation • These are electrons in the highest occupied energy level (max: 8) • Number of valence electrons can be determined by electron configuration • Inner shell electrons: not in the highest energy level Ions • An ion is a charged atom or molecule • Atoms form ions to satisfy octet rule • Will decrease to 0 or increase to 8 depending on what is easier • Ex. 2 valence electrons will lose 2 electrons to decrease to 0 • Ex. 6 valence electrons will increase by 4 to get to 8 • Cation=0 Anion=8 Atomic Structures Light • Visible Light- a kind of electromagnetic radiation, which is a form of energy that exhibits wavelike behavior as it travels through space. • All forms of electromagnetic radiation move at constant speeds of 3.0 x 10^8 m/s • Wavelength: distance between corresponding points on adjacent waves. • Frequency: the number of waves that pass a given point in a specific time Electromagnetic Spectrum Quantum Numbers • Photoelectric effect: refers to the emission of electrons from a metal when light shines on the metal • Max Planck studied the photoelectric effect, and concluded that a hot object does not continuously emit energy. Instead, it emits energy in small, specific amounts called quanta. • Quantum: the minimum number of energy that can be lost or gained by an atom. • Planck’s constant (referred to as h): 6.626 x 10^-34 J/s Quantum Models • Quantum Theory: describes mathematically the wave properties of electrons and other very small particles. • Quantum numbers: specify the properties of atomic orbitals and the properties of electrons in orbitals • The first three quantum numbers result from the Schrodinger equation, and indicate the main energy level, shape, and orientation of an orbital • Principal quantum number (n): indicates the main energy level • Values begin at n=1 continues with whole numbers • As value of n increases, the average distance from the nucleus increases • Angular quantum number (l): indicates the shape of an orbital • When L equals 0: the orbital is s, 1=p, 2=d, 3=f • S orbitals are spherical, p orbitals have dumbbell shapes, and d orbitals have clover shapes Atomic structures Electron Configuration • Electron configuration: the arrangement of electrons in an atom • First, orbital energy levels must be determined • Orbital notation: orbitals are indicated by lines, and the name written beneath. Each electron in an orbital is indicated by a half-arrow. If there are two, one must be going up and the other down. • Electron configuration notation: eliminates the lines and arrows of orbital notation. Instead, a superscript is added to the orbital name to represent the number of electrons. Examples Chart for electron configuration Rules for electron configuration • Aufbau principle: an electron occupies the lowest-energy orbital that can receive it. • Pauli exclusion principle: no two electrons in the same atom can have the same set of four quantum numbers. Thus, an orbital can hold two electrons of opposite spin. • Hund’s rule: orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin. Lewis dot diagrams • Valence electrons: the electrons available to be lost gained, or shared in the formation of chemical compounds • Since valence electrons are outer-shell electrons, their number can be determined. For example, s2p4 would have 6 valence electrons. • Electron-dot notation: an electron configuration notation in which only the valence electrons of an atom of a particular element are shown, indicated by dots placed around the element’s symbol. • Example: If you complete the electron configuration notation for fluorine, you will find that there are 7 valence electrons. In addition, you can find the element’s group on the periodic table if it is in the A or B block. Lewis dot diagrams Ions • Ion: a charged atom • Ions have either negative or positive charges through the gain or loss of electrons. • Cation: atom with a positive charge • Anion: Atom with a negative charge ATOMIC STRUTURE (WOOHOO!) BY: BITHIAH AND SHRAYA!!! WHAT IS AN ATOM? (To refresh your memory because if you don’t know by now, yikes!) 1) Atom: An atom is the basic (and smallest) unit of matter and is the “defining” structure of an element. 2) The atom was first proposed by Democritus (lucky atom!) in early 5th century BCE, where Democritus suggested that everything is composed of atoms & empty void. 3) Atoms consist of three subatomic particles, protons, neutrons, and electrons. 4) In their natural state, atoms are electrically neutral. WHAT IS THE ATOM? (CONT.) 1) 2) 3) The body of an atom is made up of two different parts, the nucleus and the electron cloud. Nucleus: Is located in the center of the atom. It consists of both protons and neutrons and is makes up most of the atom’s mass. Electron Cloud: Surrounds the nucleus and obtains most of the atom’s space. Nucleus Electron Cloud IDENTIFICATION 1) 2) Atoms are identified by their atomic number (which is the amount of protons the atom contains) An atom’s mass number is the total of protons and neutrons in the nucleus, each is worth one amu. (AMU=atomic mass unit). IONS 1) Ion: An atom with a net electric charge, due to the losing or gaining of one (or more) electrons. 2) Cations: A positively charged atom, due to losing electrons. 3) Anions: A negatively charged atom, due to gaining electrons, QUANTUM NUMBERS 1. Quantum numbers describe the location of an electron in the cloud. 2. Each electron is made up of 4 quantum numbers. 3. Every electron has a unique set of quantum numbers, like an id, zipcode, or finger print. VALENCE ELECTRONS 1. An electron located in the outer shell (valence shell) that can be shared or transferred to another atom 2. The typically reside in the highest energy level. 3. When two different atoms interact, the valence electrons determine how an atom will react in a chemical reaction by coming into contact with other atoms. ELECTRON CONFIGURATION 1) The dispersal of electrons in atomic/molecular orbitals. 2) Electrons move independently throughout the orbital. 3) From configuration, an atom's reactivity for corrosion can be determined. - Aufbau Principle: Electrons that orbit one or more atoms fill the lowest energy levels before filling the highest. - Pauli Exclusion Principle: Two electrons in the same atom can NOT have the same set of four quantum numbers. - Hund's Rule: Orbitals with the same amount of energy have only one electron before another orbital can be occupied by a second electron. All electrons in a singly occupied orbital have the same spin. 1. 2. 3. ELECTRON CONFIGURATION NOTATION Easy way to present the electrons. A subscript represents the amount of electrons in a sub-level. Ex. Neon, has 10 electrons, so its electron notation, 1s^2 2s^2 2p^6. ORBITAL NOTATION 1. Orbitals are represented by arrows with the orbital name underneath. 2. Half Arrow = One electron ATOM JOKE! Atomic Structur es Atomic Structures Atomic structure is the concept that atoms are centralized with a positively charged nucleus (contained with protons and neutrons) with negatively charged electrons surrounding it. Electromagnetic Spectrum Types of Wavelengths Longest to Shortest: Radio waves Microwaves Infrared Optical Ultraviolet X-rays Gamma Rays Wavelength and Frequency Wavelength (λ): The distance between two identical points on adjacent waves Frequency (ν) (nu): The number of waves that pass an indicated point in a given time period. Wavelength and frequency are inversely related to each other through the speed of light. c=λν Calculating Wavelength(l), Frequency(n), and Enegry (E) E= hc/λ where C=λV where V is frequency in hertz, 1/s or s-1 E is Energy in Joules (J) λ is wavelength λ is wavelength C=3.0 x 108 m/s h=6.626 x 10-34 J s C= 3.0 x 108 m/s Example Question 1. A certain microwave has a wavelength of 0.032 meters. Calculate the frequency of this microwave. Answer: 3.00 * 10^8 / .032 = 9.375 * 10^9 Atomic Structure Atomic Number: the number of protons in an atom Atomic Mass: the number of protons and neutrons in an atom Number of electrons= number of protons Valence Electrons Valence electrons are electrons in an atom, mostly in the highest energy level, that are able to gained, lost, or shared in the formation of a chemical bond. The number of valence electrons can be indicated from the period of an atom in the periodic table or its electron configuration. Maximum number of valence electrons is 8. Lewis Dot Diagram Gilbert N. Lewis created the Lewis dot diagram in 1916. The Lewis Dot Diagram is a diagram in which the valence electrons of an atom are indicated by the use of dots circling around the symbol. Example Question: 1) Draw the Lewis dot diagram for carbon: Electron Configuration The elements are represented numerically by the number of electrons in their shells and number of shells Ions Atoms or groups of atoms with a positive or negative charge. Cations: - positive charge - usually metals - fromed when an electron is taken away from an atom Anions: . - negative charge - usually nonmetals - formed when adding an electron to an atom Bibliography "A Brief Tutorial on Drawing Lewis Dot Structures." Lewisdots.html. N.p., n.d. Web. 13 Dec. 2016. <http://web.chem.ucla.edu/~harding/lewisdots.html>. "A Brief Tutorial on Drawing Lewis Dot Structures." Lewisdots.html. N.p., n.d. Web. 13 Dec. 2016. <http://web.chem.ucla.edu/~harding/lewisdots.html>. "The Definition of Valence Electron." Dictionary.com. N.p., n.d. Web. 13 Dec. 2016. <http://www.dictionary.com/browse/valenceelectron>. "The Electromagnetic Spectrum: Home Page." The Electromagnetic Spectrum: Home Page. N.p., n.d. Web. 13 Dec. 2016. <http://www.darvill.clara.net/emag/>. NASA. NASA, n.d. Web. 13 Dec. 2016. <http://imagine.gsfc.nasa.gov/science/toolbox/emspectrum1.html>. NASA. NASA, n.d. Web. 13 Dec. 2016. <http://science.hq.nasa.gov/kids/imagers/ems/waves4.html>. N.p., n.d. Web. <http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch8/>. Bibliography "Calculations between Wavelength, Frequency and Energy - Probs 1 - 10." Calculations between Wavelength, Frequency and Energy - Probs 1 - 10. N.p., n.d. Web. 13 Dec. 2016. "Quantum Numbers." Chemistry LibreTexts. N.p., 21 July 2016. Web. 13 Dec. 2016. "Wavelegnth, Frequency and Energy Calculations." Wavelegnth, Frequency and Energy Calculations. N.p., n.d. Web. 13 Dec. 2016. Atomic Structure Zainab, Julian, Sritha 3B Light and Waves • Electromagnetic radiation can be seen in terms of energy, wavelength, or frequency. • Frequency is measured in hertz • Wavelength is measured in meters • Energy is measured in electron volts. Electromagnetic Spectrum • The electromagnetic spectrum is the range of all types of EM radiation. • Radiation is energy that travels and spreads out as it goes • Types of EM radiation that make up the electromagnetic spectrum are visible light, microwaves, infrared light, ultraviolet light, Xrays and gamma-rays. Quantum Numbers • Used to describe the location of an electron within the electron cloud. • Four quantum numbers make up a complete set for each individual electron. Lewis Dot Diagram • A Lewis dot diagram is a method of writing the chemical symbol of an element by surrounding it with dots to indicate the number of valence electrons. • Valence electrons are found in an atom's outer shell and are the ones involved in chemical reactions. • Atoms tend to gain, lose or share electrons in order to acquire a full set of valence electrons, to satisfy the octet rule. Electron Configuration . • An atom's electron configuration is a numeric representation of its electron orbitals. • Eliminates lines and arrows from orbital notation. • Number of electrons in a sublevel is shown by adding to the orbital name. Valence Electrons • The valence electrons are the number of electrons an atom must lose or gain to have a full outer shell, to satisfy the Octet rule. • Highest occupied energy level Ions • Ions are charged atoms that exist when atoms have an unequal number of protons and electrons. Atomic Structures Created by Edward Vasquez and Syed Husaini Electron Configuration • Definition The organization of electrons within the orbitals of an the way the orbitals are filled out Aufbau principle: Electrons enter the lower energy orbitals to the orbitals. S,P,D, and F are all orbitals. Starting from S the lowest energy orbital highest energy orbital. Pauli Exclusion Principal: There are no two electrons that can be the Hund’s Rule: Orbitals with the same energy is singly occupied before orbital is twice as occupied.. Electrons in singly occupied orbitals amount of spin. Electron Configuration Sub- shells # of orbitals # of Electrons S 1 2 P 3 6 D 5 10 F 7 14 • Purpose: When atoms contact or interact with each other, the valence shell (outermost layer) is the part of each that will interact first. And so an atom is most reactive when it’s valence shell is not basically filled. And so an atom is most stable when it’s valence shell is full. Electron configurations can help predict various ways elements can react. Ions • Definition - An atom/molecule with an electric charge due to the loss or gain of one or more electrons. - cation: positively charged atom (lost electrons, metals) – anion: negatively charged atom (gained electrons, non metals) • Each proton carries a +1e charge • Each electron carries a -1e charge Ions - Trends • As you go down a group (periodic table), the number of electrons in the valence shell stays the same • As you go across a period (left to right) the number of electrons in the valence shell of the main group elements increases one by one. - This ignores the transition metals as they require different process, which includes the utilization of roman numerals, to determine their charge. Ions – The Formation • Atoms form ions to satisfy the octet rule • Chemical compounds form so that each atom (by gaining, losing, or sharing electrons) has an octet of electrons in its highest occupied energy level. - The number of valence electrons will either decrease to “0” or increase to 8, whichever requires less energy e.g. * an atom with 2 valence electrons will lose 2 electrons to reach 0. * an atom with 6 valence electrons will gain 2 electrons to reach 8. Lewis Dot Structure • Lewis dot structure is a quick and easy diagram that shows the valence electrons in an element. - The nucleus of the element is represented by its symbol. - The valence electrons are represented by dots placed around the symbol in pairs. • This diagram uses the Octet rule - Octet rule says that atoms like to have full outer shells of only eight electrons. - Atoms will lose or gain valence electrons to make their outer shells full with eight electrons, and they do this by bonding with other atoms. Lewis Dot Diagram(s) Lewis Dot Diagram(s) Chemical Bond(s) Valence Electrons • The electron of an atom, which is found in the valence shell. that can be transferred to or shared with another atom. - They can be transferred/shared with another atom • The number of electron shells with electrons is the same as the period number. • To find the Valence electron of an atom - Identify the group of the atom, based on the group discover the valence electrons. Valence Electrons Exp. If an element is in group 1A then it has one Valence electron. Exp. If an element is in group 1A then it has two Valence electrons. Light and Waves • Lambda (λ) : the actual distance of wavelenghts • Frequency (ν): the number of times a point on a wave passes another fixed point in one second • Equations - c=λν -E=hν - E = h c/ λ - (J) = (J·s) (m/s)/ λ Light and waves • Quantum- least amount of energy that an atom could gain or lose, a distinct packet of energy. • Photon- quantum energy that each particle carries Electromagnetic spectrum • forms of energy around us. • Deals with how charged particles have interactions with each other • Consists of wavelengths and frequencies. • Divided into seven different regions (radio waves, microwaves, infrared, visible light, ultraviolet, X-rays and gamma rays) Quantum numbers • Quantum numbers are utilized to determine and discuss the location of an electron in the electron cloud. MATTER MATTER BY: JOSE MARTINEZ AND SARRAJ AL-HILAL WHAT IS MATTER… • Matter is anything that has mass and takes up space WHAT IS CHEMISTRY? • The branch of science that deals with the identification of the substances of which matter is composed; the investigation of their properties and the ways in which they interact, combine, and change; and use of these processes to form new substances PHASES OF MATTER… SOLID • It is a definite shape • Has definite volume • Solid turns to liquid which shows it is melting • Solid turns to gas which shows it is sublimation LIQUID • It is an indefinite shape, it will take shape of any container • It has definite volume • When a liquid turns to a solid, it is freezing • When a liquid turns to a gas it is boiling then evaporates GAS • It is an indefinite shape, it will fill a container • It has indefinite volume and will fill container • When a gas turns into a liquid it is condensation • When a gas turns to solid it is deposition PLASMA • It is a high temperature physical state of matter in which atoms loose their electrons • Lighting • Fluorescent bulbs EXTENSIVE PROPERTIES A property that changes when the size of the sample changes. ex of extensive properties: mass,volume,length,and total charge. INTENSIVE PROPERTIES An Intensive property does not change when you take away some of the sample Physical Property A characteristic that can be observed or measured without changing the composition of the sample -can be described mixtures as well as pure substances Physical Change A usually reversible change in the physical properties of a substance, as size or shape Ex of physical change: Melting a sugar cube because the substance is still sugar Chemical Properties A property or characteristic of a substance that is observed during a reaction in which the chemical composition or identify of the substance is changed. Chemical Change Any change that results in the formation of new chemical substances. At the molecular level, chemical change involves making or breaking of bonds between atoms. Ex. of chemical change: Iron Rusting Classification of Matter Matter can be classified into two basic categories. Matter is either a mixture or it is a pure substance. We can classify mixtures into two categories-homogenous mixtures and heterogeneous mixtures. We can also classify pure substances into two categories- elements and compounds. Matter By: Ezra Aguilar and Joseph Kwon What is chemistry? Chemistry- The study of the composition, structure, and properties of matter and the changes it undergoes. Phases of matter Phase Changes Extensive Properties Depends on the amount of matter that is present Mass Volume Amount of energy Intensive Properties Does not depend an amount of matter that is present Melting point Boiling point Density Conductivity Physical Properties ● Characteristic can be observed/measured ○ Using the senses Examples: Color Smell Freezing point Boiling point Physical Change Change in a substance that does not change its identity Examples: phase/state Breaking apart or lessening quantity Is always reversible Chemical Properties A substance’s ability to undergo change in identity Creates a new substance How a substance reacts Combustibility Reactivity with other chemicals or substances like water or oxygen Chemical Change Change where one or more substances change into a different substance Examples: Burning Chemical reaction Cooking and baking Classification of Matter Matter can be found in two forms Pure Substance- has a fixed composition All parts have the same characteristics Can be either elements or compounds Mixture- physical blend of two or more kinds of matter where each have their own properties Homogeneous-uniform composition (kool-aid) Heterogeneous-not uniform composition Solutions are mixtures M AT T E R & C H A N G E BY: SARAH, JOSEPH, & GABBY DEFINITION REVIEW Chemistry : is the study of the structure, composition, and properties of matter and the changes it undergoes Matter: anything that has mass and takes up space Mass: a measure of the amount of matter PHASES OF MATTER SOLID definite shape & volume solid ⟶ liquid – melting solid ⟶ gas – sublimation PHASES OF MATTER LIQUID indefinite shape – takes shape of container definite volume liquid ⟶ solid – freezing liquid ⟶ gas – boiling PHASES OF MATTER GAS indefinite shape – fills container indefinite volume – fills container gas ⟶ liquid – condensation gas ⟶ solid – deposition PHASES OF MATTER PLASMA high temperature physical state of matter in which atoms lose their electrons -lightning , fluorescent bulbs PROPERTIES & CHANGES IN MATTER EXTENSIVE PROPERTIES depend on amount of matter that is present – mass – volume – amount of energy INTENSIVE PROPERTIES do not depend on amount of matter that is present – melting point – boiling point – density – conductivity PROPERTIES & CHANGES IN MATTER PHYSICAL PROPERTIES • characteristic that can be observed or measured without changing the identity of the substance • observed using the senses PHYSICAL CHANGES • a change in a substance that does not involve a change in the identity of • the substance – changes of state / phase changes – breaking apart matter, lessening quantity PROPERTIES & CHANGES IN MATTER CHEMICAL PROPERTIES relates to a substance’s ability to undergo changes that transform it into a different substance how the substance reacts – combustibility – reactivity with water, oxygen, etcetera CHEMICAL CHANGES a change in which one or more substances are converted into different substances – burning a substance – chemical reactions – cooking / baking PERIODIC TABLE PERIODIC TABLE BY: LIANN SEBASTIAN AND ALIYA SAYANI HISTORY AND DEVELOPMENT • Simple list made in 1789 by Antoine Lavoisier • 33 elements • First table published in 1869 by Dmitri Mendeleeve • Included more than 60 elements • Ordered by increasing atomic weight • Modern periodic table created in 1945 • Included 118 elements ORGANIZATION AND LAYOUT • Modern periodic table: • Arranged by atomic mass across rows • Arranged by physical and chemical properties • Horizontal rows= periods, numbered 1-7 • Columns= groups, numbered 1-18 • Can be split into A and B blocks • • • Metals TYPES OF ELEMENTS • Left side of the periodic table • Solid at room temperature • Silver-gray luster • Malleable and ductile Nonmetals • Located on right side • Solid or gas at room temperature (except bromine) • Poor conductor of heat and electricity • Brittle, no common color Metalloids • Has characteristics of both metals and nonmetals • Between metals and nonmetals • Semi-conductors • Boron, silicon, germanium, arsenic, antimony, tellurium CHEMICAL FAMILIES • • • • Alkali Metals • Group 1 or 1A • Most reactive, especially with nonmetals Alkaline Earth Metals • Group 2 or 2A • Reactive, especially with water Transition Metals • Group 3 or 3A • Less reactive than alkali/alkaline earth metals Rare Earth Metals • Lanthanide series- period 6 below table • Actinide series- period 7 below table CHEMICAL FAMILIES CONT. • Other Metals • Group 13- aluminum, gallium, indium, thallium • Group 14- tin, lead • Group 15- bismuth • Group 16-polonium • Halogens • Group 17 or 7A • Most reactive nonmetals, especially with metals • Noble gases • Group 18 or 8A • Unreactive, colorless, odorless PERIODIC TRENDS • Atomic Radius- measure of the size of an elements atoms • Across a period: decreases • Down a group: increases • Ionic Radius- distance between the nucleus and the electron in the outermost shell of an ion • Metals: always cations • Nonmetals: always anions • Ionization Energy- amount of energy needed to remove the valence electron to form a cation • Across a period: increases • Down a group: decreases • Electron Affinity- a neutral atom’s likelihood of gaining an electron • Across a period: decreases • Down a group: increases • Electronegativity- measure of the ability of an atom to attract electrons • Across a period: increases • Down a group: decreases NOBLE GAS NOTATION • Shorthand for electron configuration • Beginning of electron configuration is represented by the symbol of the noble gas in the period above, in brackets • Remaining portion of electron configuration is normal for desired element • • Ex: • Ne: 1s22s22p6 • Al: 1s22s22p63s23p1 Noble Gas Notation= Al:[Ne] 3s23p1 The Periodic Table By: Jacob Knackstedt and Austin McCrae Antoine Lavoisier First official periodic table published. Antoine described elements as a material which could not be decomposed into a simpler form. Contained 33 elements out of our know 118. It was made in to columns. It was organized by metals, gases, earths, and nonmetals Dimitri Mendeleev • • • • • Second published periodic table. It was published in the year 1869. Contained 60 plus elements it looked like a jigsaw puzzle. Elements were ordered by their atomic weight. Set in columns. Modern periodic table • • • • • • Created in 1945 Total of 118 known elements Rows are arranged by atomic mass Columns arranged by elements properties Horizontal rows 1-7 called periods. Columns called groups organized 1-18. Separated as a and b block. Element types: Metals • • • • • • Good conductors of both heat and electricity. Most are solid at a room temp. Silver and gray-white in appearance. Malleable Ductile On the left side of the periodic table Alkali metals • • • • • • 1a elements Silver Relatively soft Most reactive out of all elements Not found in nature as a free element Only 1 valence electron Alkaline earth metals • • • • 2a elements Silver colored Reactive, but nit as reactive as Alkali Stronger than Alkali metals. Transition metals • • • • B block elements High luster value Not as reactive as Alkali or Alkaline Mercury exists as liquid in room temp. Rare earth metals • • • • • • 2 types Lanthanides Period 6 and below Actinides Period 7 and below All are radioactive elements Other metals • • • • 3a: aluminum, gallium, thallium, indium 4a: lead, tin 5a:bismuth 6a:polonium Nonmetals • • • • Not a good conductor of electricity and heat. Can be gases or solid at room temperature. Very delicate Don’t share a common color between all of them Halogens • • • • • • • 7a Are highly reactive with metals Most reactive out of the nonmetals Liquids in this group are Bromine. Gases in this group are Chlorine and Fluorine. Solids in the group are Astatine and Iodine. Are toxic elements. Noble gases • • • • • 8a Gases Most of the time are unreactive Odorless and colorless Maximum amount od electrons in valence shell Metalloids • • • • They have some of the attributes of both metals and nonmetals Found in between the metals and the nonmetals Semiconductors Less malleable than the metal elements Periodic trends Noble gas notation • When doing electron configuration notation write the noble gas from the period above then continue writing notation. • Phosphorus: • Chlorine: [Ne]3s 2 3p 3 [Ne]3s 2 3p 5 Bibliography • • • • • • • http://www.rsc.org/education/teachers/resources/periodictable/pre16/develop/mendeleev.htm http://www.chemicalelements.com/groups/alkali.html http://www.chemicalelements.com/groups/rareearth.html http://www.chemicalelements.com/groups/nonmetals.html http://www.chemicalelements.com/groups/noblegases.html file:///C:/Users/Jacob%20Knackstedt/Downloads/Periodic_Table_of_the_Elements.pdf Notes Periodic Table By: Puja Nayak & Rawan Layth HISTORY AND DEVELOPMENT • Dmitiri Mendeleev published the first periodic table in 1869 • • • • • More than 60 elements Rotated Order by increasing atomic weight New columns with repeating characteristics Modern table made in 1945 Development • • • • • 118 elements Arranged by atomic mass across rows Arranged by physical and chemical qualities in columns Horizontal rows (periods) 1-7 Vertical columns (families) 1-18 A block and B block About the Table TYPES OF ELEMENTS • an element that is a good conductor of heat and electricity • left “side” of the periodic table • solid at room temperature • most have silver or gray-white luster • malleable: can be hammered or rolled into thin sheets • ductile: can be drawn into a thin wire METALS ▪ group 1 or 1A ▪ silver in appearance ▪ soft enough to cut with a knife ▪ react vigorously with nonmetals ALKALI METALS ▪ group 2 or 2A ▪ harder, denser, stronger than alkali metals ▪ silver colored ▪ react with water ▪ reactive metals ALKALI EARTH METALS ▪ group 3 – group 12 elements – B block elements ▪ good conductors ▪ high luster ▪ less reactive ▪ mercury exists as a liquid at room temperature TRANSITION METALS ▪ lanthanide series – period 6 below table – most are found in the Earth’s crust ▪ actinide series – period 7 below table ▪ these elements sometimes categorized as transition metals RARE EARTH METALS ▪ from group 13 or 3A elements – aluminum, gallium, indium, thallium ▪ from group 14 or 4A elements – tin, lead ▪ from group 15 or 5A elements – bismuth ▪ from group 16 or 6A elements – polonium OTHER METALS NONMETALS ▪ an element that is a poor conductor of heat and electricity ▪ right “side” of the periodic table ▪ either solid or gas at room temperature exception of bromine (liquid at room temperature) ▪ brittle ▪ no common color NONMETALS ▪ group 17 or 7A elements ▪ most reactive nonmetals ▪ toxic HALOGENS ▪ group 18 or 8A elements ▪ gases at room temperature ▪ odorless and colorless ▪ generally unreactive NOBLE GASES ▪ an element that has some characteristics of metals and some characteristics of nonmetals ▪ found between the metals and nonmetals (“stair step”) ▪ solid at room temperature ▪ less malleable than metals, not as brittle as nonmetals ▪ semiconductors ▪ boron, silicon, germanium, arsenic, antimony, tellurium METALLOID Periodic Trends • Shorter way to write electron configuration • Beginning of electron configuration is represented by a symbol of the noble gas in the period above the wanted element. • The rest is written as normal • Example: Al: [Ne]3s 23p 1 NOBLE GAS NOTATION ORBITAL • # of electrons (in a-block) can be determined by its location on the periodic table • Al found in group 3A= 3 valence electrons VALENCE ELECTRONS • Can be determined from valence electrons • Add or remove to satisfy octet rule Valence Electrons Electrons gained/lost Ion charge 1 1 lost 1+ 2 2 lost 2+ 3 3 lost 3+ 4 4 lost/4 gained 4+ 4- 5 3 gained 3- 6 2 gained 2- 7 1 gained 1- 8 none none IONIC CHARGE • Half distance of nuclei of atoms bonded together Down a group Across a period Increasing positive charge of nucleus ATOMIC RADIUS Caused when electrons go to higher energy levels further away from the nucleus. Anion Cation Removing valence electrons=smaller electron cloud IONIC RADIUS Electron cloud spread out because electrons not drawn towards nucleus • Energy needed to remove an electron from the atom (neutral) • Each ionization level increases Across a period Down a group Metals lose electrons easily, nonmetals do not IONIZATION ENERGY Electrons are further from nucleus so easier to remove • Energy change when an electron is gained (Neutral atom) Across a period Energy released by positive value NOT FAVORABLE Down a group Energy released by negative value FAVORABLE Nonmetals easily gain electrons while metals don’t ELECTRON AFFINITY Hard to add electrons to large atoms • The ability for an atom to attract electrons (chemical compounds) Down a group Across a period Larger radius so electrons feel less pull ELECTRONEGATIVITY Electrons attracted to nonmetals PERIODIC TABLE BY: MELLENY , QAYLAH, AHAN HISTORY AND DEVELOPMENT • ANTOINE LAVOISIER- made the first list of 33 elements • JOHANN W. DOBEREINER- found 3 elemts with similar characteristics - Lithium, sodium, and potassium made one triad • JOHN NEWLANDS- arranged all elments known at the time according to incressing amotic mass • LOTHAR MEYER- arranged 56 elemts into one table according to their characteristics • DMITRI MENDELEEVGrouped elements by similar chemical properties and arranged elements by incressing mass • HENRY MOSELY- arranged elements by increasing proton number ORGNIZATION AND LAYOUT • 18 vertical columns called families or groups • 7 horizontal rows called periods or series • Split into A nad B block • Arranged by atomic mass across rows TYPES OF ELEMENTS • Metals – good conducter of heat and electricity left side of the periodic table solid at room temp. • Non Metals – poor conductors of heat and electricity On the right side of the periodic table Solids or gasses at room temp. • Metalloids- some characteristics of nonmetals and others of metals Found between metals and nonmetals (stair step) Solid at room temp. CHEMICAL FAMILIES • ALKALI METALS ( Group 1) Most reactive metals • ALKALINE EARTH METALS (Group 2) Less reactive than group one • TRANSITION METALS ( Group 3-12) Good conductors of electricity • Halogens( Group 17) Most reactive nonmetals • Noble Gases (Group 18) Rarely react with other elements because they have a full valence shell PERIODIC TRENDS • Atomic Radius- size of the atom Increases down and left • Ionic Radius- Adding an electron makes atom bigger Taking away an electron makes it smaller • Ionization Energy – energy requred to remove an electron from the atom increases up and right • Electron Affinity- how much an atom wants to gain Increases up and right • Electronegativity- ability of atom to hold electrons tightly Increases up and right NOBEL GAS NOTATION • Short cut method for electron arrangement • EXAMPLE:: • Na 1s^2 2s^2 2p^6 3s^1 • -Use Neon • - Use neon’s configuration as [Ne] • - Use in Na: [Ne] 3s^1 COVALENT COMPOUND STRUCTURES Covalent Compound Structures LEWIS STRUCTURES, VSEPR THEORY, AND INTERMOLECULAR FORCES CAROLINA ROMERO MADELEINE ROOK Lewis Structure Formula which represents a nucleus and inner shell electrons using covalent bonds to show unshared electrons Covalent Bond: The sharing of an electron pair between 2 atoms that result in a chemical bond. Lone bond. Pair: A pair (2) electrons not involved in a Lewis Structures Step 1 Ex: PCl3 Count the number of valence electrons in each of the elements in the compound -1 Phosphorus atom: 5 valence electrons 1(5)=5 é -3 Chlorine atoms: 7 valence electrons 3(7)=21 é Add the numbers of totaled electrons - 21é+5é=26é Lewis Structures Step 2 Put the least electronegative element in the center, then all the other elements surrounding it. Connect the outer elements to the center element. Ex: PCl3 Subtract original é from the number on bonds used Cl P Cl Cl each bond is 2 electrons 26é-6é=20é Lewis Structures Step 3 Fill all the outside elements with the rest of the electrons. *One bond is 2é 20é Cl P Cl Cl Lewis Structures Exceptions More Than: -Some atoms have more than 8 -Non-metals in period 3 or below -Add a lone pair to central atom if electrons remain Ex: NO2-1 O 14é- 2é N O Less Than: -less than 8 -duet (2): Hydrogen -quartet (4): Beryllium -sextet (6): Boron Ex: CH3Cl 6é H Cl C H H VSEPR Theory Stands for Valence Shell Electron Pair Repulsion 3 dimensional arrangements of all domains (EDG) 3 dimensional arrangement of bonded pairs (MDG) Angle between bonded pair domains (Bond Angle) Intermolecular Forces forces between molecules (covalently bonded compound) London Dispersion Model (weakest) - found in every molecule (constant motion of electrons) Intermolecular Forces Dipole-Dipole -Created by equal or opposite charges with short distances of separation + I - Cl Cl I+ Intermolecular Forces Hydrogen Bonding (strongest) -Highly electronegative atom bonded with Hydrogen. -Creates a very strong dipole + H H - H+ O H + O - MOLECULAR STRUCTURES BY AKUL, JAYA, JOSH VSEPR THEORY Valence Shell Electron Pair Repulsion VSEPR theory is a model used to predict the geometry of specific molecules by the number of electron pairs around the central atom The basic idea of VSEPR theory is that the valance electrons will repel each other and go into a specific geometric formation that reduces the amount of repulsion that they undergo VSEPR THEORY Electron Domain Geometry The number of lone pairs and bonded pairs around a particular atom in a molecule. Molecular Domain The number of lone pairs around a particular atom in a molecule. Geometry INTERMOLECULAR FORCES Forces of attraction between molecules (covalently bonded compounds) London-dispersion forces Dipole-Dipole Hydrogen bonding LONDON-DISPERSION FORCES Weakest of the three forces -Cause by constant motion of electrons Everything has Londondispersion Forces DIPOLE DIPOLE Attracted forces between the positive end and the negative end of a polar molecule Polar and Non Polar NO DIPOLE Carbon - Hydrogen bond The same element HYDROGEN BONDING Strongest of the three forces -results from hydrogens bonded to a very electronegative atom -creates strong dipole -no permanent attraction There are three ways a hydrogen atom can be formed -Have fon Fluorine Oxygen Nitrogen QUIZ 1. What does it mean to have FON? 2. Name two intermolecular forces? 3. What is VSEPR theory? CHEMICAL FORMULAS AND NAMES CHEMICAL FORMULAS AND NAMES SOHAN GADE & MARIJ AHMED 3A CHEMISTRY OVERVIEW • Ionic Bonds / compounds • Covalent bonds / compounds • Polyatomic ions • Nomenclature COVALENT BONDS • Bonds in which electrons are shared between atoms of different elements • These are formed with two non-metals • Example: H2O, CO2 • The least electronegative element is named first • To name the compound you take the first element, look at how many atoms there are of that compound, and put an appropriate prefix (e.g. tetra for four), if it is one then there is no prefix • Take the second element, and do the same process and add the prefix, and then add the suffix, -ide IONIC BONDS • Ionic Bonds are formed when two elements either give up or gain electrons. • Metals will form cations(lose electrons) and nonmetals will form anions (gain electrons). • This differs from covalent bonding, because the atoms lose or gain the electrons, rather than sharing them • The atoms involved will usually try to follow the octet rule (having a full/stable eight electrons in the outer shell) • Examples : NaCl (Table Salt), 𝐶𝑎𝐶𝑙2 (Calcium Chloride) POLYATOMIC IONS • Binary- Has only two elements, ex. NaCl • Polyatomic- Multiple atoms with collective charge (e.g. Ammonium, Hydroxide, Cyanide) • Base unit has suffix of –ate, -ite has one less Oxygen, -per has one more, and hypohas two less than the base • For some, it is the Hydrogen that changes, -bi means one more Hydrogen, and –di means two more Hydrogen • If there are multiple of the same polyatomic ion, the formula of the ion is put in parentheses and a subscript indicating the number used is placed outside the parentheses WITH TRANSITION METALS • Find which charge of the metal that needs to be used to make the compound neutral • Take the element name (cation) place it first, and then take the charge and rewrite it a Roman numeral, then place the other element/polyatomic ion name • Example: Copper (II) Sulfate NAMING IONIC COMPOUNDS • If it is a nonmetal and metal that are bonding, the ending of the second element is changed to –ide, the cation comes first. • With transition metals, the transition metal cation’s Roman numeral is used after the element name to indicate the charge, followed by the second element with the suffix –ide, ex. MgO (Magnesium Oxide) • With polyatomic ions, the cation is placed first and then polyatomic ion name is placed (ex. Copper (II) Hydroxide) Chemical Formulas and Names By: Emanuel Encarnacion & Michael Rudolph Ionic Bonds/Compounds Ionic: formed when neutral atom gains or loses an electron. Cation: formed when a neutral atom loses an electron and becomes positively charged. Anion: formed when a neutral atom gains an electron and becomes negatively charged. Ionic Compound: a compound formed from a cation and an anion. Typically metal + nonmetal reaction Ex. Potassium Nitride K^1+ N^3- K^1+K^1+K^1=K^3 K3N3 balanced Covalent Bonds: Bonds between atoms when they share electrons Covalent Bonds/ Compounds Covalent Compounds: Formed when nonmetal atoms share valence electrons Polyatomic Ions Polyatomic ions: a group of atoms that have a charge. Polyatomic ions usually have a charge because the collection of atoms has either gained an extra electron or else it has lost an electron. Nomenclature: Choosing names for things, especially in science Ionic Compounds: Metals with only one charge are written as their basic name; metals with multiple charges are written with roman numerals with specific charge Nomenclature Fe₃O₂ - iron(II) oxide Covalent Compounds: First element has element name with prefix of numer of atoms present; second element has prefix of number of atoms present and an ending of -ide CHEMICAL FORMULAS AND NAMES BY: Kabir Bhakta & Kaushik Jampala 2B Chemistry POLYATOMIC IONS What is a polyatomic ion? • A polyatomic ion is a charged group of covalently bonded atoms. Therefore, they are either cations or anions created by an imbalance between the amount of electrons and protons • The prefix poly- means many, so a polyatomic ion is an ion with more than one atom. • Note that these are named differently than monoatomic ions which contain only one atom. POLYATOMIC IONS Naming Polyatomic Ions • The following are rules for identifying different polyatomic ions based on the number of oxygens they have. They are identified using special prefixes and suffixes. • • • • The suffix-ite is added to the ion with one less oxygen atom. The suffix –ate is added to the most common ion. The prefix hypo- is placed in front of the ion with one less oxygen than the –ite ion The prefix per- is placed in front of the ion with one more oxygen than the –ate ion. NAMING POLYATOMIC ION COMPOUNDS • The following are the rules for naming ionic compounds containing polyatomic ions: 1. The name of the cation is written first. 2. This is followed by the name of the anion. 3. However, when there is multiple of a polyatomic ion present, parenthesis are placed around the entire ion. IONIC BONDS • Ionic bonding is a type of chemical bond that involves two, oppositely charged ions • Ionic bonding occurs when one atom takes an electron from another atom hence the phrase “taken, not shared” • The metal loses electrons to become positively charged and the nonmetal receives those electrons to become a negatively charged anion IONIC COMPOUNDS When naming ionic compounds follow these rules: • An ionic compound is named first by its cation and then by its anion • The cation keeps the same name as its element and the anion name’s ending is replaced with –ide • Ex: NaF = sodium fluoride • If the cation is a transition metal, the charge is written in Roman numerals in parentheses directly after the element name • Ex: Cu(NO3)2 = copper (II) nitrate • If either the cation or anion is a polyatomic ion, the polyatomic ion name is used in the overall compound, the polyatomic ion always stays the same COVALENT BONDS • Covalent bonds form between two nonmetals • Unlike ionic bonding, covalent bonding “shares” electron pairs between atoms • Covalent bonding is generally weaker than ionic bonding • Ex: H2O = Oxygen needs two more valence electrons to fill its outer shell, and hydrogen needs to get rid of one electron to satisfy the octet rule so both hydrogen atoms will share their extra electron with oxygen creating a balanced compound COVALENT COMPOUNDS When naming covalent compounds follow the rules below: • The first element when naming will be the less electronegative element • A prefix will be added to the first element unless its –mono • The second element in the compound will be named using the appropriate prefix and changing the ending to –ide • Ex: SO2 = sulfur dioxide Chemical Formulas and Names Ashton Mathew Noah Kim Ionic Bonds/Compounds Ionic compound formula = Cation + Anion Cation - positively charged ion Anion - negatively charged ion Cation naming 1. Group A – element name 2. Transition metals- element name and the charge (with Roman Numerals) 3. Polyatomic – ion name Anion naming 1. Nonmetal - element name + “-ide” 2. Polyatomic – ion name Covalent Bonds/Compounds Covalent Bonding = Results from the sharing of electron pairs between two atoms 2 elements are used to form a covalent bond First element + Second element = covalent compound First Element 1. Write the element name 2. Add a prefix if there are more than one of the element (never use mono) Prefixes 1-mono 2-di 3-tri 4-tetra 5-penta 6-hexa 7-hepta 8-octa 9-nona 10-deca Second Element 1. Write the element name with the ending “-ide” 2. Add a prefix in the front if there are more than one of the element (mono can be used) Polyatomic Ions Polyatomic Ion: Molecular ion that is a charged chemical species composed of two or more atoms covalently bonded or of a metal complex that can be considered to be acting as a single unit. Prefixes “ate”: original ion “ite”: subtract one oxygen “per”: add one oxygen “hypo”: adds two oxygen or subtracts two oxygen based on ending “ite” or “ate” List Acetate: CH3O2(-) Ammonium:NH4(+) Bromate:BrO3(-) Carbonate:CO3(2-) Chlorate:CIO3(-) Chromate:CrO4(2-) Cyanide:CN(-) Dichromate:Cr2O7(2-) Iodate: IO3(-) Hydroxide: OH(-) Nitrate:NO3(-) Oxalate:C2O4(2-) Phosphate:PO4(3-) Sulfate:SO4(2-) MOLES AND CHEMICAL FORMULAS MOLES AND CHEMICAL FORMULAS By: Aftab Shaik, Aaron Purewal, Madar Marar Mole Concept • A mole is a unit of measurement like a dozen is a unit of measurement that means 12 of a certain thing • In simple terms, it is the amount of substance that contains the same amount of particles in 12 grams of carbon-12 • According to the National Institute of Standards and Technology (the NIST) the formal definition of 1 mole is a: • The mole is the amount of a substance of a system which contains as many elementary entities as there are atoms in 0.012 of carbon-12; its symbol is "mol." When the mole is used, the elementary entities must be specified and may be atoms, molecules, ions, electrons, other particles, or specified groups of such particles. Exactly how big is this number? • Avogadro’s number: the amount of particles in exactly one mole of a pure substance • This number is represented as 6.022*10^23 • This number is used as a conversion factor to convert between particles/atoms and moles for the same chemical • Moles of a substance to atoms of that substance: multiply the mole amount by Avogadro’s number. You will end up with atoms of that substance. • Atoms of a substance to moles of that substance: divide the atom amount by Avogadro’s number. You will end up with moles of that substance. Molar Mass • This is as the mass of 1 mol of the substance • • Measured in grams per mol (g/mol) This measurement can be found by taking the avg. atomic mass of all the elements in the substance • You have to take the avg. atomic mass of each individual element and add all of these together and you will result in the molar mass of the substance Molar Mass as a Conversion Factor • You can use the molar mass of a substance to convert between moles and grams of that substance • Moles of a substance to grams of that substance: Can be found by taking the mol amount and multiplying the molar mass of that substance and you will end up with grams of that substance • Grams of a substance to moles of that substance: Can be found by taking the grams amount and dividing this amount by the molar mass of the substance and you will end up with moles of that substance Percent Composition • The percent composition is the percent by mass of each element in a compound • This can be found by taking the mass of a specific element and dividing that by the total mass of the compound • • This will result in a number less than 1 in decimal form We then take this decimal and multiply it by 100% to get the percent composition Empirical Formula • The empirical formula of a compound is the smallest whole number ratio between a compound • • Ionic compound’s ratios must always be reduced The steps in order to find this are as follows: 1. 2. 3. 4. 5. You will be given a measurement, either grams or percent composition of the compound Find the grams of each part of the compound Divide this measurement by the molar mass of each individual chemical This new measurement will be in moles and divide these numbers by the smallest number of moles (out of all the individual chemicals) in the result This will end with a small ratio that can be made into a formula Molecular Formula • The molecular formula of a compound is the ratio between elements in one molecule of compound • In order to find this, we must first find the empirical formula of the compound • Then, we will most likely be given the molar mass of the compound • Using the molar mass of the empirical formula, we must divide the given molar mass by the molar mass we found with the empirical formula and we will find a conversion factor • We take this number and multiply each of the element amounts of the empirical formula and you will end up with the molecular formula MOLES AND CHEMICAL FORMULAS BY: NEEL AND MOSTAFA MOLE CONCEPT • MOLE - AMOUNT OF SUBSTANCE THAT CONTAINS AS MANY PARTICLES AS THERE ARE ARE ATOMS IN EXACTLY 12 GRAMS OF CARBON-12 • AVAGADRO’S NUMBER – NUMBER OF PARTICLES IN ONE MOLE OF A SUBSTANCE • AVAGADRO’S CONSTANT = 6.022 X 10^23 PERCENT COMPOSITION • MULTIPLY EACH AMOUNT OF ELEMENTS BY THEIR OWN MOLAR MASS • DIVIDE EACH BY THE TOTAL MOLAR MASS OF THE COMPOUND • MULTIPLY 100 MOLAR MASS • MOLAR MASS = ATOMIC MASS OF AN ELEMENT • MOLAR MASS OF COMPOUND = SUM OF ALL MOLAR MASS IN THE COMPOUND EMPIRICAL FORMULA • ASSUME 100 GRAMS OF ALL ELEMENTS • DIVIDE BY MOLAR MASS OF THAT ELEMENT • DIVIDE BY SMALLER VALUE • ROUND ALL NUMBERS • THESE NUMBERS WILL BE THE SUBSCRIPT OF THE ELEMENTS MOLECULAR FORMULA • DIVIDE GIVEN AMOUNT OF GRAMS OF COMPOUND BY MOLAR MASS OF COMPOUND • ROUND THE PRODUCT • MULTIPLY PRODUCT BY EACH SUBSCRIPT IN THE EMPIRICAL FORMULA MOLES & CHEMICAL FORMULAS Mole Concept • A unit of a substance that is equivalent to the amount of atoms as there are in 12 grams of carbon-12 – Represented by Avogadro’s Number, or NA – 6.022*1023 • Avogadro’s number can be used as a conversion factor when calculating the number of particles or atoms within an amount of a substance • 1 mole of a substance=6.022*1023 atoms of that substance • Atoms to moles: divide by Avogadro’s number • Moles to atoms: multiply by Avogadro’s number Mole Concept (cont.) • Example: Determine the number of moles in 6.755 atoms of silicon. • Calculation: (6.755 atoms silicon)*(1 mol silicon/6.022*1023 atoms silicon)= 1.122*10-23 mol silicon • Example: Determine the number of atoms in 0.00384 moles of francium. • Calculation: (o.oo384 mol francium)*(6.022*1023 atoms francium/1 mol francium)=2.05 atoms francium Your Turn! • Determine the number of moles in 3.95 atoms of potassium – 6.56*10-24 mol K • Determine the number of atoms in 2.56 moles of nitrogen – 1.54*1024 Molar Mass • The mass of one mole of a substance • Measured in units of grams per mole, or g/mol • Molar mass of an element is its average atomic mass – Ex: magnesium’s molar mass is 22.990 g/mol • The molar mass of a compound is the combined average atomic masses of each of the elements – Ex: LiF: 6.941+18.998= 25.939 g/mol • You try! Determine the molar masses of the following compounds: – NaCl: • 58.443 g/mol – H2O: • 18.015 g/mol – MgCl2: • 95.211 g/mol Molar Mass as a Conversion Factor • Applicable when performing dimensional analysis • Allows conversions between grams and moles of the same element • Grams to moles: divide by molar mass • Moles to grams: multiply by molar mass • Example: determine how many grams are present in 2.14 moles of carbon dioxide • Calculation: (2.14 mol CO2)*(44.009 g CO2/1 mol CO2) =94.2 g CO2 • Example: determine how many moles are present in 96.74 grams of potassium chloride • Calculation: (96.74 g KCl)*(1 mol KCl/74.551g KCl)=1.298 mol KCl You Try! • Determine the number of moles in 16.54 grams of H2O – 0.9181 mol H2O • Determine the number of grams in 5.20 moles of calcium chloride – 577 g CaCL2 Percent Composition • Percentage by mass of each element in a compound • Divide mass of one element by the compound’s molar mass and multiply by 100 • Example: FrF • (223/241.998)*100= 92.1% Fr – Subtract 92.1 from 100 to determine percent composition of fluoride – 7.9% • You try! Find the percent compositions of each of the elements in CH2O – C: 40.001% – H: 6.714% – O: 53.284% • Percentages should add up to 100 Empirical & Molecular Formulas • Empirical Formula: The smallest whole number ratio between elements in a compound – NaCl, H2O are empirical formulas – Ionic compounds are always written in empirical formulas • Molecular formula: ratio between elements as they actually exist in one molecule of a covalent compound – C8H18 is a molecular formula • Empirical formulas can be derived from molecular formulas – C4H9 is the empirical formula for C8H18 Finding Empirical Formulas from Percent Compositio ns Determine the empirical formula of a compound that contains 52.11% carbon, 13.14% hydrogen, and 34.75% oxygen. 1. Assume there are 100 grams of the compound. – 52.11g C, 13.14g H, 34.75g O 2. Divide by respective molar masses to determine number of moles of each element – 52.11g C/12.011= 4.33852 mol C – 13.14g H/1.008= 13.03571 mol H – 34.75g O/15.999=2.17201 mol O 3. Divide these values by the smallest value to determine a whole number ratio. Round. – 4.33852 mol C/2.17201=~2 mol C – 13.03571 mol H/2.17201=~6 mol H – 2.17201 mol O/2.17201= 2 mol O Finding Empirical Formulas from Percent Compositio ns 4. Write out empirical formula using the determined number of moles of each element. – C2H6O2 • Molecular formulas require a given molar mass to be determined. • Given: 186.204 g/mol is the molar mass. Find the molecular formula of C2H6O2 1. First, find the empirical formula mass – 2(12.011)+6(1.008)+2(15.999)=62.068 g/mol 2. Then , divide the given molar mass by the empirical formula mass. – 186.204/62.068= 3 3. Multiply the subscripts by this value to determine the molecular formula. – C6H18O6 Your turn! 1. Determine the empirical formula of a compound with 63.50% silver, 8.25% nitrogen, and 28.25% oxygen. – AgNO3 2. The given molar mass is 339.78 g/mol. Find the molecular formula. – Ag2N2O6 Moles and Chemical Formulas By: Ihsan and Amirtha Mole Concept – Mole: The amount of a substance that contains as many particles as there are atoms in 12 grams of carbon-12 – Avagadro’s Number: Used as a conversion factor to connect between particles and goals for the same chemical – 6.022 x 1023 𝑚𝑜𝑙1− – Converting from moles to atoms: multiply the amount of moles by Avagadro’s number, 6.022 x 1023 𝑚𝑜𝑙1− – Converting from atoms to moles: divide the amount of atoms by Avagadro’s number, 6.022 x 1023 𝑚𝑜𝑙1− Examples 1) 8.08 x 1023 atoms Kr 8.08 x 1023 atoms Kr ? mol Kr 2) 1.04 mol Au ? atoms Au 1.04 mol Au / x 6.022 x 1023 atoms Kr 6.022 x 1023 atoms Au = 1.34 mol Kr = 6.26 x 1023 atoms Au Molar Mass – Molar Mass: The mass of one mole of a substance – Units of grams per mole (g/mol) – Molar Mass of an element is equal to the average atomic mass of the element – Molar mass of a compound is equal to the sum of the masses of the elements present Examples 1) 𝑀𝑔𝐶𝑙2 2) 𝐶13 𝐻18 𝑂2 24.305+ (2 x 34.453) = 95.211 g/mol (12.011 x 13) + ( 1.008 x 18) + (15.999 x 2) = 206.285 g/mol Percent Composition – Percent by mass for each element in a compound – Mass of the element/ total mass of compound x 100= % of that element in the compound 1) Calculate the mass of element A 2) Calculate the mass of element B 3) Divide the mass of element A by the total mass of the compound and multiply by 100 4) Divide the mass of element B by the total mass of the compound and multiply by 100 Examples 1) 𝐻2 𝑂 2) 𝐹𝑒𝑆 H: (1.008 x 2) = 2.016 g Fe: 55.845 g 18.015 g 87.911 g O: ( 1 x 15.999) = 15.999 g S: 32.066 g (2.016/18.015) x 100 = 11.19 % H ( 55.845/87.911) x 100 = 63.52 % Fe (15.999/18.015) x 100 = 88.81 % O ( 32.066/87.911) x 100 = 36.48 % S Empirical/Molecular Formulas – Molecular Formula: The ratio between elements represents the number of atoms in one molecule of the molecular compound – Empirical Formula: An empirical formula consists of the symbols for the elements combined in a compound, with subscripts showing the smallest whole number mole ratio of the different atoms in the compound *Ionic compounds only have empirical formulas *Covalent compounds may have both Deriving Empirical and Molecular Formulas Empirical Formula Molecular Formula – If percent composition is given, first convert into grams (assume 100 g) – A molar mass is always given – Divide by molar mass – Take the smallest mole amount and divide the other mole amounts by it – Round to the closest whole number and write out the empirical formula based on the ratios – Determine the molar mass of the empirical formula derived – If the 2 amounts are different, divide the molar mass by the empirical formula mass – Round the quotient to nearest whole # – Multiply the empirical formula by that value to derive the molecular formula Example: Empirical Formula Given: 63.52% Fe, 36.48% S Fe- 63.52 g (molar mass 55.485 g/mol) Assume 100g S- 36.48 g (molar mass 32.066 g/mol) 63.52/ 55.485 = 1.13743 mol Fe/ 1.13743 ≈ 1 36.48/ 32.066 = 1.13765 mol S / 1.13765 ≈ 1 Ratio- 1:1, so FeS Empirical Molecular Given: 𝑃2 𝑂5 Empirical molar mass- 283.89 g/mol Molecular formula-? Actual molar mass- 141.943 g/mol 283.89/141.943= 2.00002818≈ 2 So,(𝑃2 𝑂5 )2 𝑃4 𝑂10 Molecular Formula: 𝑃4 𝑂10 Moles & Chemical Formulas By: Ryan Kwon & Ashwin Thevarumuriyil The Mole • Mole: an amount of a substance • Avogadro’s Number: the number of particles in one mole of a pure substance ▫ 6.022 x 10^23 mol^-1 ▫ Used as a conversion factor to convert particles (atoms to moles) Molar Mass • Molar Mass: The mass of one mole of a substance ▫ Measured in grams per mole g/mol ▫ Used as conversion factor to convert between grams and moles • Molar mass of an element is equal to atomic mass ▫ Oxygen has molar mass of 15.999 g/mol • Molar mass of compound equal to sum of individual molar masses ▫ CO^2 = (1 x 12.011) + (2 x 15.999) = 44.009 g/mol Percent Composition • Percent Composition: Percent of mass of each element in a compound. ▫ (Part/whole) x 100 = % ▫ (Mass of element/molar mass) x 100 = % H202 H: (2 x 1.008) = 2.016 O: (2 x 15.999) = 31.998 ▫ Molar Mass = 34. 014 ▫ H: 5.93% ▫ O: 94.07 Empirical & Molecular Formulas • Empirical Formula: Formula with the smallest whole number ratio between elements in a compound ▫ Ionic compounds are always empirical MgCl2 • Molecular Formula: Formula with ratio between elements as they exist in one molecule of a molecular compound. ▫ Can be empirical formula H20 How To Find Empirical Formula • 63.52% Fe 36.48% S 1. Assume 100g for example 2. Divide by molar mass ▫ ▫ 63.52% Fe -> 63.52g Fe/55.845 = 1.13743 mol Fe 36.48% S -> 36.48g S/32.066 = 1.13765 mol S 3. Divide by smallest number of moles ▫ ▫ 1.13743/1.13743 = 1 mol Fe 1.13765/1.13743 = 1.00019 mol S 4. Round Special Cases to Finding Empirical Formula • After dividing by smallest number of moles, the quotient may end in one of the following ▫ .25 (x4) ▫ .33 (x3) ▫ .5 (x2) ▫ .67 (x3) ▫ .75 (x4) • If you end up with these numbers, just multiply by numbers next to them until they become close enough to a whole number to round REACTIONS Balancing equations __Al +__HCl ->__AlCl3 +__H2 Balancing equations _2_Al +_6_HCl ->_2_AlCl3 +_3_H2 combustion • When a chemical includes hydrogen and carbon as reactants and always produces carbon dioxide and water as a product _1_CH4+_2_O2->_1_CO2+_2_H2O YOU WANT SPICY? HERE WE GO Process into Equation What happened between A and B? What about C and D? Omg really are they still dating? Long story short?? RUMORS TEXT TALK (aka equation); AB + C -> AC + B ..get ready to learn you some more! it ain't over yet yung'uns Process into Words What happened between A and B? What about C and D? Omg really are they still dating? Tell me in full detail?? RUMORS LONG STORY (aka Words); AyyBee and Cee yielded AyyCee and Bee Synthesis When two or more substances combine to form a new product. You find the one.. They're cute, they give you butterflies, and most of all they're good hearted. Y'all think it's real and you'll last forever. Awh, good for you. #synthesis A+B -> AB Decomposition When a compound breaks down to form two or more elements, basically the opposite of synthesis You remember that time you and that special person thought y'all were gonna last forever? Lolyeahmetoo #decomposition AB->A+B Double Replacement When to compounds react to form two new compounds by switching elements You're with your bae, and your best friend is with his bae, now your best friend has your ex bae, and you have your best friend's ex bae. . #ol'switcharoo AB+CD->AD+CB Single Replacement When an element reacts with a compound and replaces an element within the compound Lol nice bae you got there…Don't mind if I do… AB+C->AC+B REACTIONS By: Dhruv Kota and Hasta Zain HOW TO TRANSLATE EQUATIONS TO WORDS? • Translating chemical equations to words is done by using number of moles and type of chemical. • The format for writing a chemical equation is “__ moles of Chemical and __ moles of chemical yields __ moles of chemical and __ moles of chemical. • Example- 4NH3+ 5O2 4NO+6H2O • Answer- four moles ammonia and five moles of oxygen gas yield four moles of nitrogen monoxide and six moles of water. TRANSLATING EQUATIONS TO FORMULAS • To translate a chemical equation to the formula, the same process as above must be used, but just reversed. • The worded equation must be used, and the correct amount of moles and chemical names are written for the chemical reaction formula. • Example- One mole of sodium chloride and one mole of sulfuric acid yield one mole of hydrochloric acid and one mole of sodium hydrogen sulfate. • Answer- 1NaCl + 1H2SO4 1HCl + 1NaHSO4 TYPES OF REACTIONS • There are 5 types of reactions- Synthesis, Decomposition, single replacement, double replacement and combustion. • Synthesis: When the reactants all combine into one product. • Decomposition: When the reactant splits into different products. • Single replacement: When an element in the reactants take the place of another element so the composition of the reactants end up switched with one element in the product. • Double replacement: When 2 compounds react, the positive ions in the reactant react with the negative ions in the reactant switch places with each other causing them to be switched in the product. • Combustion: Reactions that occur when something burns-uses organic molecules. EXAMPLES OF TYPES OF REACTIONS • Synthesis: Mg + O2 ---> MgO • Decomposition: HgO ---> Hg + O2 • Combustion: CH4 + O2 ---> CO2 + H2O • Single replacement: CH4 + O2 ---> CO2 + H2O • Double replacement: KOH + H2SO4 ---> K2SO4 + H2O BALANCING EQUATIONS • Chemical equations are written to express different chemical reactions. • Most of them need to be balanced-they need to have the same amount of each atoms on each side of the equation. • How to balance chemical equations: 1. Write equation. 2. Figure out a coefficient to put next to each compound that would make each side have equal amounts of each element. • Ex: • C3H8 + 5O2 --> 4H2O + 3CO2 STOICHIOMETRY Stoichiometry CHIRAYU T. AND KIRK A. How to Convert Converting grams, moles, and atoms Basically dimensional analysis Converting Grams Grams to grams Grams to moles 20g C = ?mol C Grams to atoms 20g C = ?atoms C Converting Moles Moles to moles Moles to grams 1mol Na = ?g Na Moles to atoms 1mol Na = ?atoms Na Converting Atoms Atoms to atoms Atoms to moles 5×1020 atoms K = ?mol K Atoms to grams 20×1016 atoms Cl = ?g Cl Converting Between Chemicals Uses ratios between chemicals From the reaction: B2H6 + O2 HBO2 + H2O a. What mass of O2 will be needed to burn 36.1 g of B2H6? b. How many moles of water are produced from 19.2 g of B2H6? STOICHIOMETRY REVIEW By: Aleena Hassan and Raisha Ahmed CONVERTING MOLES TO MOLES Mole ratio (Equation) Amount of GIVEN substance (in mol) mol UNKNOWN ——————— mol GIVEN Given in the problem Conversion factor Amount of UNKNOWN substance (in mol) Calculated CONVERTING MOLES TO MOLES EXAMPLE If you have 2.00 mol of N2 reacting with H2, how many moles of NH3 will be produced? Equation: N2 + 3H2 2NH3 2.00 mol N2 Given in the problem 2 mol of NH3 ——————— 1 mol N2 Conversion factor (mole ratio) 4.00 mol NH3 Calculated CONVERTING MOLES TO GRAMS Mole ratio (Equation) Amount of GIVEN substance (in mol) Given in the problem mol UNKNOWN ——————— mol GIVEN Molar mass (Periodic table) molar mass of UNKNOWN (in g/mol) Conversion factors Mass of UNKNOWN substance (in g) Calculated CONVERTING MOLES TO GRAMS EXAMPLE If you have 3.00 mol of H2O react with CO2, how many grams of C6H1206 are produced? Equation: 6CO2+6H20 C6H1206+6C02 Mole ratio (Equation) 3.00 mol H2O Given in the problem 1 mol C6H1206 ——————— 6 mol H2O Molar mass (Periodic table) 180.18 g C6H1206 ———————— 1 mol C6H1206 Conversion factors 90.1 g C6H1206 Calculated CONVERTING GRAMS TO MOLES Inverted molar mass (Periodic table) Mass of GIVEN substance (in g) Given in the problem 1 ————————————— molar mass of GIVEN (in g/mol) Conversion factors Mole ratio (Equation) mol UNKNOWN ——————— mol GIVEN Amount of UNKNOWN substance (in mol) Calculated CONVERTING GRAMS TO MOLES EXAMPLE How many moles of mercury (II) oxide, HgO, are needed to produce 125 g of oxygen, O2? Equation: 2Hg0 2Hg + O2 Inverted molar mass (Periodic table) 125 g O2 Given in the problem 1 mol O2 ————— 31.998 g O2 Mole ratio (Equation) 2 mol HgO ——————— 1 mol O2 Conversion factors 7.81 mol HgO Calculated CONVERTING GRAMS TO GRAMS Inverted molar mass (Periodic table) Mass of GIVEN substance (in g) Given in the problem Mole ratio (Equation) 1 ————————————— molar mass of GIVEN (in g/mol) mol UNKNOWN ——————— mol GIVEN Conversion factors Molar mass (Periodic table) molar mass of UNKNOWN (in g/mol) Mass of UNKNOWN substance (in g) Calculated CONVERTING GRAMS TO GRAMS EXAMPLE You are given this equation: 4Al +3O2 2Al2O3 If you were given 74.00 grams of Al, how many grams of Al2O3 would be produced? Inverted molar mass (Periodic table) 74.00 g Al Given in the problem Mole ratio (Equation) 1 mol Al ————————————— 26.982 g Al 2 mol Al2O3—— ———— 4 mol Al Conversion factors Molar mass (Periodic table) 101.96 g Al2O3— ————— 1 mol Al2O3 139.8 grams of Al2O3 Calculated CONVERTING MOLES TO ATOMS Amount of element (in mol) Given in the problem 6.022*1023 Number of atoms of element Calculated CONVERTING MOLES TO ATOMS EXAMPLE If you have 2.00 mol of Helium, how many atoms do you have? 2.00 mol He Given in the problem 6.022*1023 1.20*1024 atoms He Calculated CONVERTING ATOMS TO MOLES Number of atoms of element Given in the problem 1 ————— 6.022*1023 Amount of element (in mol) Calculated CONVERTING ATOMS TO MOLES EXAMPLE If you have 5 atoms of Sulfur, how many moles of Sulfur do you have? 5 S atoms Given in the problem 1 mol S ————— 6.022*1023 S atoms 8 mol S Calculated CONVERTING GRAMS TO ATOMS Mass of element (in grams) Given in the problem 1 ————————— molar mass of element 1 ————— 6.022*1023 Number of atoms of element Calculated CONVERTING GRAMS TO ATOMS EXAMPLE How many atoms of Sulfur are in 4.00 g of Sulfur? 4.00 g S Given in the problem 1 mol S ————————— 32.065 g S 1 atom S ————— 6.022*1023 mol S 7.51*10-22 atoms S Calculated CONVERTING ATOMS TO GRAMS Number of atoms of element Given in the problem 1 ————— 6.022*1023 Molar mass of element Mass of element (in grams) Calculated CONVERTING ATOMS TO GRAMS EXAMPLE If you have 1 atom of Hydrogen, how many grams of Hydrogen do you have? 1 atom H Given in the problem 1 mol H ————— 6.022*1023 atoms H 1.008 g H —————— 1 mol H 2gH Calculated DETERMINING LIMITING REACTANT Limiting reactant: the reactant that limits the amounts of other reactants that can combine and the amount of product that can form in a chemical equation To determine: Balance the equation, convert to moles, use stoichiometry for each individual reactant to find the mass of product produced. The reactant that produces a lesser amount of product is the limiting reactant. DETERMINING EXCESS REACTANT Excess reactant: the substance that is not used up completely in a reaction To determine: Balance the equation, convert to moles, use stoichiometry for each individual reactant to find the mass of product produced. The reactant that produces a larger amount of product is the excess reactant. CALCULATING THEORETICAL YIELD Theoretical yield: the maximum amount of product that can be produced from a given amount of reactant In most chemical reactions, the amount of the product obtained is less than the theoretical yield To calculate: find the moles in the limiting reactant, and convert it from moles to grams CALCULATING PERCENT YIELD Percent yield: the ratio of the actual yield to the theoretical yield , multiplied by 100 Actual yield: the measured amount of a product obtained from a reaction Example: If the actual yield is 20 and the theoretical yield is 25, the percent yield is 80%