* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
Download Types of Measurement
Electrolysis of water wikipedia , lookup
Physical organic chemistry wikipedia , lookup
Nuclear binding energy wikipedia , lookup
Metastable inner-shell molecular state wikipedia , lookup
X-ray photoelectron spectroscopy wikipedia , lookup
Stoichiometry wikipedia , lookup
X-ray fluorescence wikipedia , lookup
Hypervalent molecule wikipedia , lookup
Condensed matter physics wikipedia , lookup
Electronegativity wikipedia , lookup
Gas chromatography–mass spectrometry wikipedia , lookup
Resonance (chemistry) wikipedia , lookup
Nuclear transmutation wikipedia , lookup
Particle-size distribution wikipedia , lookup
Abundance of the chemical elements wikipedia , lookup
Metallic bonding wikipedia , lookup
Chemical element wikipedia , lookup
Chemical thermodynamics wikipedia , lookup
Elementary particle wikipedia , lookup
Electron configuration wikipedia , lookup
Molecular dynamics wikipedia , lookup
Rutherford backscattering spectrometry wikipedia , lookup
IUPAC nomenclature of inorganic chemistry 2005 wikipedia , lookup
Chemical bond wikipedia , lookup
Atomic nucleus wikipedia , lookup
Periodic table wikipedia , lookup
History of molecular theory wikipedia , lookup
History of chemistry wikipedia , lookup
State of matter wikipedia , lookup
Extended periodic table wikipedia , lookup
GHSGT Chemistry Review Types of Measurement Qualitative measurement: uses words to describe something. (i.e. A yellow-green gas was released.) Quantitative measurement: uses numbers to describe something. (i.e. The oxide has a mass of 1.567 grams.) Temperature Scales Celsius, Kelvin, Fahrenheit Converting Celsius to Kelvin: C + 273 = ____ K a. Example: 25 C = ______ 25 + 273 = 298 K Converting Kelvin to Celsius: K – 273 = ______ C a. Example: 300 K = _________ C 300 – 273 = 27 C Classification of Matter The law of conservation of matter states that matter can neither be created nor destroyed. anything that has mass Matter and occupies space Are the properties and composition constant? no mixture physical change seawater mud fruit salad yes yes pure substance Is chemical separation into simpler substances possible? no chemical compound element water, sodium chloride oxygen, gold, sulfur Classification of Matter Compound: A substance with a constant composition that can be broken down into elements by chemical processes. Element: A substance that cannot be broken down into simpler substances by chemical means. Classification of Matter Mixture 1. Components of two or more components that can be separated by physical means, such as chromatography, filtration, or distillation 2. No chemical bonds between the separate parts 3. Components can vary in composition. Classification of Matter Two kinds of mixtures 1. Homogeneous mixtures (solutions) A. can contain solids, liquids, or gases B. are evenly mixed, appear as one component even though there are two or more parts. 1. Examples: plain jello, air, and salt water Classification of Matter 2. Heterogeneous mixtures A. are unevenly mixed B. can see separate components C. Examples: dirt, concrete, salad dressing Classification of Matter Classify these as an element, a compound, a heterogeneous mixture, or a homogeneous mixture. 1. Concrete (heterogeneous mixture) 2. Air (homogeneous mixture) 3. Salt (compound) 4. Gold (element) 5. Helium (element) 6. Tea (homogeneous mixture) 7. Salt water (homogeneous mixture) Metric Prefixes Metric Prefixes and Powers of 10 1. mega = 106 (megagrams) (Mg) 2. kilo = 103 (kilograms) (kg) 3. base = 100 (grams) (g) 4. centi = 10-2 (centimeters) (cm) 5. milli = 10-3 (milligrams) (mg) 6. micro = 10-6 (micrograms) (μg) 6000 mg = 6 g = 0.006 kg Precision vs. Accuracy Precision refers to how close a series of measurements are to one another. 1. Replication of result / value Accuracy refers to how close a measured value is to an accepted value. 1. True or correct result / value Seven Base SI Units 1. Length 2. Mass 3. Time 4. Temperature 5. Current 6. Amount 7. Luminousity meter (m) kilogram (kg) second (s) Kelvin (K) Ampere (amp) mole (mol) candela (cd) Derived Units Derived Units are combinations of two or more units. 1. Examples: speed = mi/hr, ft/s, m/s density = mass/volume area = (length) (width) Density Density is the mass of a substance per unit volume of the substance. 1. density = mass/volume (g/ml or g/cm3) 2. Water (the standard for all density values) = 1.0 g/ml 3. 1ml = 1cm3 4. What is the volume of a 50 g metal block with a density of 5 g/cm3 ? 5. d = m/V, m = dV, V = m/d The Chemists’ Shorthand: Atomic Symbols mass number = atomic number = 39 K 19 = element symbol Atomic Structure atomic number = the number of protons 1. The atomic number also represents number of electrons in an neutral atom mass number = the total number of protons plus neutrons in the nucleus Atomic Structure In uncharged (neutral) atoms, the atomic number equals the number of protons equals the number of electrons. If an atom is charged (different number of protons and electrons), then it is called an ion. 1. A (+) charged ion is called a cation. 2. A (–) charged ion is called an anion. Atomic Structure Isotopes are atoms of the same element (with the same number of protons) but different number of neutrons. 1. An element is defined by the number of protons in the atom’s nucleus. A. Example: carbon – 12 (12 neutrons); carbon – 14 (14 neutrons) Periodic Table of the Elements Atomic Structure Atomic Structure Atomic Structure Particle Mass Charge Electron 9.1094 x 10-31 kg -1 Proton 1.6726 x 10-27 kg +1 Neutron 1.6749 x 10-27 kg 0 Periodic Table of the Elements Groups (vertical - up and down the table) 1A = alkali metals 2A = alkaline earth metals 7A = halogens 8A = noble gases Periods (horizontal - across the table) numbered 1-7 Periodic Table of the Elements Periodic Table of the Elements Metals vs Nonmetals - see staircase on the Periodic table 1. Metals - to the left of the staircase; mostly solids; conduct electricity; lose electrons to form positive ions 2. Nonmetals - to the right of the staircase; most are gases; nonconductors of electricity; gain electrons to form negative ions 3. Metalloids - border the staircase; have properties of metals and nonmetals. Periodic Table of the Elements Periodic Table of the Elements There are 7 periods on the periodic table numbered 1 - 7. 1. They represent the major energy levels (n). 2. They are horizontal rows that extend from left to right. Ex: Period 2 includes Li - Ne. Periodic Table of the Elements Periodic Table of the Elements Groups 1. IA - Alkali Metals (1 valence electron) very reactive 2. IIA - Alkaline Earth Metals (2 valence electrons) 3. VIIA - Halogens (7 valence electrons) very reactive 4. VIIIA – Noble Gases (8 valence electrons except for Helium) non-reactive (very stable) Periodic Table of the Elements Periodic Table of the Elements Representative Elements - The Group “A” Elements which include all the Groups IA to VIIIA Transition Elements - The Group “B” Elements Periodic Table of the Elements Periodic Table of the Elements Lanthanide Series - the 4f row that includes # 57 (Lanthanum) through # 71 Lu Actinide Series - the 5f row that includes # 89 Ac (Actinum) through # 102 No Periodic Table of the Elements Periodic Table of the Elements Atomic Radius - the distance from the center of the nucleus to the outermost valence shell Periodic Trend - The atomic radius increases as one moves down the group. The atomic radius decreases as one moves across a period. Periodic Table of the Elements Inorganic Compound Classification Two main kinds of compounds 1. Ionic: made up of ions of opposite charge A. strong electrostatic force of attraction; ionic bond B. electrons are transferred 2. Covalent: made up of two or more nonmetals A. electrons are shared Chemical Bonds Atoms bond with each other to become more chemically stable than they were before they bonded. To do this their outer electron (valence) shell must be complete. 1. The Octet Rule states that atoms will either gain, lose , or share valence electrons to attain “8” electrons in their outer (valence) shell to become stable. 2. The Noble Gases are the only one group of elements that are already stable. Naming Binary Ionic Compounds Ions of opposite charge are bonded together. 1. The metal cation (+) is written first and is named by the metal’s name 2. The nonmetal anion (-) is written second and is named by the nonmetal’s name with a revised ending of - ide. 3. Net charge of ions in the compound = 0. 4. Subscripts are used to indicate the number of ions needed to attain the necessary net charge of 0. Naming Binary Ionic Compounds Examples of Binary Ionic Compounds What would the formulas be? 1. Sodium chloride 2. Lithium nitride 3. Calcium Fluoride What would the names be? 1. Al2S3 2. BaO 3. MgBr2 Periodic Table of the Elements Covalent (Molecular) Compounds Made up of nonmetals that share electrons between atoms. 1. This type of bond is called a covalent bond. Naming Binary Covalent Compounds 1. The first nonmetal’s name is that of the element. 2. The second nonmetal’s name has an -ide ending, just like with ionic compounds. 3. Use prefixes to describe the subscripts (1 - mono; 2 - di; 3 - tri; 4 - tetra; 5 - penta; 6 - hexa; 7- hepta; 8 - octa; 9 - nona; 10 - deca) Examples of Covalent Compounds P2O5 - diphosphorus pentoxide How would you write the following formulas? 1. Carbon monoxide 2. Tetranitrogen decoxide 3. Dinitrogen hexoxide Balancing Chemical Equations The Law of Conservation of Matter 1. The mass of the products = the mass of the reactants (matter is not created or destroyed). 2. Chemical equations are balanced to ensure that a chemical reaction follows the law of conservation of matter. Five Balancing Equation Guidelines 1. Count the number of atoms of each element on both the reactant and the product side. 2. Use coefficients (the numbers in front of the chemical symbol or formula) 3. Never add or change the subscripts. 4. There are Seven Diatomic Elements (N2, O2, F2, Cl2, Br2, I2, H2) 5. Balance the hydrogen atoms and the oxygen atoms last. Balancing Chemical Equations 6. A chemical equation uses symbolic language to describe a chemical reaction. 7. Equation means equal numbers of atoms of each element on both sides. 8. Quantities of reactants and products are expressed in moles by using coefficients. Balancing Chemical Equations Symbols for balancing equations (s) – solid; (g) – gas; (l) – liquid (aq) – aqueous (dissolved in water) Go to www.usaprep.com to practice balancing equations. Balancing Chemical Equations Try these equations: 1. H2 + O2 H 2O 2. Ca + Br2 3. Ba(s) + 4. Mg + CaBr2 O2(g) AuCl3 BaO(s) MgCl2 + Au Balancing Chemical Equations The Answers: 1. 2H2 + O2 2H2O 2. Ca + Br2 CaBr2 3. 2Ba(s) + O2(g) 2BaO(s) 4. 3Mg + 2AuCl3 3MgCl2 + 2Au Acids and Bases Acid Definitions: 1. Sour taste 2. Neutralize the actions of bases 3. Blue litmus paper turns red 4. Liberates hydrogen gas when reacted with certain metals 5. Examples: Foods and Drinks Acids and Bases Base Definitions: 1. Bitter taste 2. Neutralizes the action of acids 3. Slippery to the touch 4. Red litmus paper turns blue 5. Examples: Cleaning Solutions Acids and Bases an acid + a base a salt + water Strong acids: hydrochloric acid, sulfuric acid, nitric acid Strong base: sodium hydroxide Weak acid: acetic acid Weak base: ammonia pH Scale The pH scale has values from 0 - 14. 0 - 6 is acidic 7 is neutral 8 -14 is basic The Mole The mole is the SI unit of measure for the amount of a substance. 1. The mole is a way to measure the mass of elements and compounds. 2. The molar mass of any element is numerically equal to its atomic mass and has the units of g/mol. 3. Example: The mass of one mole of potassium (K) is 40 grams. Physical and Chemical Changes 1. Physical Change - Only appearance changes. Identity still the same. Kinds of changes are: Size/Shape and Phase (State). A. Size and Shape (splitting, breaking, tearing, hammering, etc.). B. Phase (State) (melting, vaporizing, freezing, condensing, etc.). Physical and Chemical Changes 2. Chemical Change - The appearance and the identity changes. A new product is formed. The way you know a new product is formed is by the following: A. gas formed B. color change C. mass change D. heat change E. solid formed F. light released Examples are: rusting, growing, burning, cooking, combusting, fermenting, frying, and exploding. Nuclear Reactions A nuclear reaction results when an unstable nucleus breaks down and emits radioactive particles. 1. There are 3 types of particles released during radioactive decay. A. alpha particle (α) released during alpha decay B. beta particle (β) released during beta decay C. gamma ray (γ) released during any decay Three Types of Radioactive Decay 1. alpha particle (helium nucleus (+ 2 charge) A. largest, slowest, least penetrating particle 2. beta particle (fast electron (– 1 charge) A. more penetrating than an alpha particle 3. Gamma radiation (no particle, no mass or charge – pure energy) A. most penetrating and most damaging (shielded by lead) Half Life The half life of a radioactive isotope is the time it takes for one half of the isotope to decay. 1. Example: The half life of mercury - 195 is 31 hours. If you start with 20 g, how much will be left after (A) 31 hours? (10 g) (B) 62 hours? (5 g) Particles in an Electric Field Remember the charges for each particle. Also remember that like charges repel and opposite charges attract. Nuclear Fission and Nuclear Fusion Nuclear Fission is the splitting of large nuclei resulting in a tremendous release of energy. Nuclear Fusion is the combining of small nuclei resulting in an even greater release of energy. The sun uses nuclear fusion to produce energy. 1. 2 hydrogen atoms combine to form 1 helium atom. Energy/ Heat/ Phase Changes 1. Temperature – the measure of the average kinetic energy of particles. 2. Heat – form of energy that may be absorbed or released. Flows from a warm body to a cooler one until equilibrium is reached. 3. A calorie – the amount of heat required to raise the temperature of one g of water one degree Celsius. Types of Phase Changes 1. melting - solid to liquid 2. freezing - liquid to solid 3. evaporation - liquid to gas 4. condensation - gas to liquid 5. sublimation - solid to gas 6. deposition - gas to solid Key Terms for Heat 1. Energy – the capacity to do work 2. Heat – energy transferred from one object to another (The SI unit of heat is the Joule (4.18 Joules = 1 calorie). 3. Thermochemistry – the study of heat effects in chemical reactions 4. Combustion – chemical reactions that release heat 5. Exothermic reaction – one that releases heat 6. Endothermic reaction – one that absorbs heat Heat Specific heat – amount of heat required to raise the temperature of one gram of a substance 1o Celsius. Specific heat of water is 1. 1. Water – high specific heat due to hydrogen bonds. Water doesn’t change temperature very much despite large amounts of heat energy added or subtracted. 2. Metals have low specific heat capacities. Energy in Chemical Reactions The bond breaking that occurs in reactants during a chemical reaction requires energy. The bond formation that occurs in products during a chemical reaction releases energy. An endothermic or exothermic reaction is determined by the balance between these two processes. 3 Phases (States) of Matter The balance between the attractive forces between the particles and the kinetic energy of the particles determines the phase of matter. 1. High KE of the particles and low attractive forces between the particles = gas. 2. Low KE of the particles and high attractive forces between the particles = solid. 3. Intermediate KE of the particles and intermediate attractive forces between the particles = liquid. The Phase Diagram for Water Solids The two main types of solids are crystalline solids and amorphous solids. 1. In a crystalline solid, the atoms, ions, or molecules are arranged in an orderly, repeating, 3-dimensional pattern (crystal lattice). Solids 2. In an amorphous solid, the internal structure lacks order. Atoms, ions, or molecules are randomly arranged. A. These substances usually cool rapidly – Not enough time for the particles to arrange themselves in a pattern. B. Examples: Rubber, glass, plastics, polymers Allotropes Allotropes – substances with same elemental composition, different geometric arrangements. 1. Example, carbon has 4 allotropes: A. diamond - formed under tremendous pressure B. graphite - more loosely packed C. soot - randomly bonded (amorphous form) D. buckey ball Kinetic Molecular Theory of Gases 1. Gases – made up of very small particles (atoms or molecules) and separated by large distances (low density). 2. Particles are in constant, random, straight-line motion – thousands of perfectly elastic collisions per second. Total KE remains constant. 3. Gases exert a pressure due to the collisions on each other. Collision Theory 1. When gas particles collide they exert a pressure on their container. 2. Temperature is the measure of the average kinetic energy of the gas particles. 3. There are four main properties of gases that determine their physical behavior. Four Main Properties of Gases 1. Pressure = force / area 2. Temperature = the average kinetic energy of particles 3. Volume = space occupied by matter. 4. Amount of gas is measured in grams or moles. What happens to the volume of a gas if the pressure is increased at constant temperature? What happens to the volume of a gas if the temperature is increased at constant pressure? Lewis Dot Structures Rules for Lewis Dot Structures 1. Count total number of dots (valence electrons) in the structure. 2. Spatially arrange the atoms. (More than two atoms – find the central atom) 3. Obtain 8 dots around each atom (2 around hydrogen). 4. If single bonds don’t work, try double bonds, then triple bonds. Three Factors Affecting the Rate of Dissolving How fast will a solute dissolve in a solvent? 1. Stirring or agitation (more solute / solvent contact at a faster rate) 2. Smaller particles (increases the surface area of the solute, therefore there is more solute / solvent contact at a faster rate) 3. Increasing the temperature (increases the kinetic energy and faster rate of contact between the solute/solvent particles) The Solubility of Solids at Different Temperatures 1. Solubility of sugar at 50oC? 2. Which solute least affected by increase in temp? 3. Generally, as temp increases, solubility? GHSGT Chemistry Review The End