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Midterm Review Unit 1 Matter and Measure Unit 2 Atomic Theory and Structure Unit 3 Nuclear Chemistry Unit 4 Periodic Table Unit 5 Bonding Unit 6 Naming and Moles Unit 1 - Matter & Measure Chapters 1-3 Matter Anything that has mass and takes up space, volume Classified into two categories ◦ Substances (Pure) ◦ Mixtures Atom Simplest form of matter Made up of Subatomic Particles Different atoms have different properties Pure Substances Element ◦ simplest form of matter that has a unique set of properties. ◦ Can’t be broken down by chemical means Compounds ◦ substance of two or more elements chemically combined in a fixed proportion ◦ Can be broken down by chemical means Mixtures Physical blend of two or more substances Homogeneous ◦ Composition is uniform throughout ◦ Solution is a homogeneous mixture ◦ Aqueous Solution is something mixed in water Heterogeneous ◦ Composition is not uniform throughout 2.3 Matter Separating Mixtures Differences in physical properties can be used to separate mixtures Examples: ◦ Filtration – Separates solids from liquids in heterogeneous mixtures ◦ Distillation – Separates homogeneous liquid mixtures based on different boiling points ◦ Evaporation – evaporate away liquid to leave solid ◦ Chromatography – separation of substances based on polarity Phases(States) of Matter Solid (s) ◦ Definite shape and volume ◦ Particles are packed tightly together Liquid (l) ◦ Definite volume, takes shape of container ◦ Particles can slide past each other Gas (g) ◦ Takes shape and volume of container ◦ Particles are spread very far apart Phase Changes Solid Liquid Liquid Solid Liquid Gas Gas Liquid Solid Gas Gas Solid Melting Freezing Vaporization Condensation Sublimation Deposition Temperature does NOT change during a phase change Aqueous Solutions Dissolved in water (aq) used after chemical symbols ◦ NaCl(aq) Temperature Measure related to the heat of an object Measured in °Celsius or Kelvin(no degrees) Conversion K C 273 Density Amount of matter in a given amount of space Amount of mass in a given volume m D V Identifying Substances Physical Property ◦ quality or condition of a substance that can be observed or measured without changing the substance’s composition ◦ Ex: Color, shape, size, mass Physical Change ◦ some properties change, but the composition remains the same ◦ Ex: melting, freezing, tearing Identifying Substances (cont) Chemical Change ◦ change that produces matter with a different composition than the original matter ◦ Ex. burning, rusting, decomposing, exploding, corroding Chemical property ◦ property that can only be observed by changing the composition of the substance. ◦ Ex: Reactivity with acids, reactivity with oxygen Energy Ability to do something Exothermic ◦ Process when energy is released or given off ◦ Ex: Burning, freezing Endothermic ◦ Process when energy is absorbed or taken in ◦ Ex: Melting Scientific Method Observation - make observations Hypothesis - proposed explanation Experiment - test Analyze Data - check to see if results support hypothesis. Theory - well tested explanation Law - concise statement that summarizes the results Scientific Laws Law of Conservation of Mass ◦ Mass can not be created or destroyed, only changed into different forms Law of Conservation of Energy ◦ Energy can not be created or destroyed, only changed into different forms Metric System Prefixes Prefix Power Symbol Giga 109 G Mega 106 M kilo 103 k deci 10-1 d centi 10-2 c milli 10-3 m micro 10-6 μ nano 10-9 n pico 10-12 p Significant Figures If the decimal point is present, start counting digits from the Pacific (left) side, starting with the first non-zero digit. 123 0.00310 (3 sig. figs.) If the decimal point is absent, start counting digits from the Atlantic (right) side, starting with the first non-zero digit. 32 1 31,400 (3 sig. figs.) SigFigs for Math Addition and Subtraction ◦ Answer has to have the same number of decimal places as least number of decimal places in what you are adding or subtracting Multiplication and Division ◦ Answer has to have same number of Sigfigs as least number of Sigfigs in what you are multiplying or dividing Percent Error Main Menu Atomic Theory and Structure Chapters 4-5 Atomic Theories Democritus ~ 400 BC ◦ believed that atoms were indivisible and indestructible Dalton ~ 1800’s ◦ Developed through experiments ◦ First Atomic Model Dalton’s Atomic Model All elements are composed of tiny indivisible particles called atoms Atoms of the same element are identical. The atoms of any one element are different from those of any other element. Dalton’s Atomic Model (cont) Atoms of different elements can physically mix together or can chemically combine in simple whole-number ratios to form compounds. Chemical reactions occur when atoms are separated, joined, or rearranged. Atoms of one element, however, are never changed into atoms of another element as a result of a chemical reaction. “Plum Pudding” Model ◦ Discovery of electron requires new model Uniform positive sphere with negatively charged electrons embedded within. Rutherford Gold Foil Experiment Shot alpha particles at gold foil Most went through foil with little or no deflection. Some were deflected at large angle and some straight back. Rutherford Model Conclusions from Gold Foil Experiment ◦ Atom is Mostly Empty Space ◦ Dense positive nucleus ◦ Electrons moving randomly around nucleus Bohr Model Dense positive nucleus Electrons in specified circular paths, called energy levels These energy levels gave results in agreement with experiments for the hydrogen atom. Bohr Model Bohr Model Each energy level can only hold up to a certain number of electrons Level 1 2 electrons Level 2 8 electrons Level 3 18 electrons Level 4 32 electrons Wave Mechanical Model More detailed view of the Bohr Model Still maintains nucleus with protons and neutrons Electrons are found in orbitals within the energy levels Regions of space where there is a high probability of finding an electron Modern Model ◦ AKA Quantum Mechanical Model, Electron Cloud Model Subatomic Particles Electron ◦ Discovered first ◦ Negative charge (-1) ◦ Approx mass ~ 0u Proton ◦ Discovered second ◦ Positive charge (+1) ◦ Approx mass ~ 1u Neutron ◦ Discovered last ◦ No charge (0) ◦ Approx mass ~ 1u Just slightly larger than a proton Atomic Structure Atomic Number ◦ Number of protons in an element ◦ All atoms of the same element have the same number of protons Mass Number ◦ Number of protons and neutrons in an atom Atomic Structure # of Neutrons = Mass Number – Atomic Number Atoms of the same elements can have different numbers of neutrons Isotope – atoms of the same element with different number of neutrons Ion Atom or group of atoms that have gained or lost one or more electrons ◦ Have a charge Example: ◦ H+, Ca2+, Cl-, OH- Average Atomic Mass Atomic Mass ◦ Weighted average based on the relative abundance and mass number for all naturally occurring isotopes Example ◦ C-12 98.9% 12.011u ◦ C-13 1.1% ◦ 0.989*12 + 0.011*13 = 12.011u Electron Configuration The way in which electrons are arranged in the atom Example: Na 2-8-1 Valence Electrons ◦ Electrons in the outermost energy level Ground State vs. Excited State Ground State ◦ When the electrons are in the lowest available energy level ◦ Ex: Na 2-8-1 Excited State ◦ When one or more electrons are not in the lowest available energy level ◦ Ex: Na 2-7-2 or 2-8-0-1 or 2-6-1-1-1 Energy Level Transitions Gaining energy will move an electron outward to a higher energy level (Absorption) When an electron falls inward to a lower energy level, it releases a certain amount of energy as light (Emission) Line Spectra Emission Spectra ◦ Shows only the light that is emitted from an electron transition Absorption Spectra ◦ Shows a continuous color with certain wavelengths of light missing (absorbed) Main Menu Unit 3 - Nuclear Chemistry Chapter 25 Radioisotopes Nuclei of unstable isotopes are called radioisotopes. An unstable nucleus releases energy and/or mass by emitting radiation during the process of radioactive decay Radiation Type Alpha Mass 4 Charge +2 Strength Weakest Symbol α Beta Gamma 0 0 -1 0 middle Strongest β γ Nuclear Stability For smaller atoms, a ratio of 1:1 neutrons to protons helps to maintain stability ◦ C-12, N-14, O-16 For larger atoms, more neutrons than protons are required to maintain stability ◦ Pb-207, Au-198, Ta-181 Nuclear Reactions Unstable isotopes of one element are transformed into stable isotopes of a different element. They are not affected by outside factors, like temp and pressure. They can not be sped up or slowed down. Radioactive Decay Radioisotopes will undergo decay reactions to become more stable Alpha Decay Beta Decay Positron Emission Alpha Decay 220 87 Fr At 4 2 216 85 Beta Decay 90 38 Sr Y 0 1 90 39 Positron Emission 19 10 Ne e F 0 1 19 9 Transmutations Any reaction where one element is transformed into a different element Natural ◦ Usually has one reactant ◦ Alpha and Beta Decay Artificial ◦ Usually has more than one reactant ◦ Fission, Fusion Example 9 4 Be H He Li 1 1 4 2 6 X 3 Fission Splitting of a larger atom into two or more smaller pieces ◦ Nuclear Power Plants One Example: U n Ba Kr 3 n 235 92 1 0 141 56 92 36 1 0 Fusion Joining of two or more smaller pieces to make a larger piece ◦ Sun, Stars Examples: 4 H He 2 2 2 3 1 1 H 1 H 2 He 0 n 2 1 1 1 4 2 0 1 H H He n 3 1 4 2 1 0 Energy Production Energy is produced by a small amount of mass being converted to energy ◦ Fission and Fusion ◦ More energy is produced by fusion than any other source E=mc2 Fission vs. Fusion Advantages of Fission ◦ Produces a lot of energy ◦ Can be a controlled reaction ◦ Material is somewhat abundant Disadvantages of Fission ◦ Uses hazardous material ◦ Produces hazardous material Long Half Life ◦ Reaction can run out of control. ◦ Limited amount of fissionable material Fission vs. Fusion Advantages of Fusion ◦ Lighter weight material ◦ Easily available material ◦ Produces waste that is lighter and has shorter half-life ◦ Produces more energy than fission Disadvantages of Fusion ◦ Must be done at very high temperatures Only been able to attain 3,000,000K ◦ Have not been able to sustain stable reaction for energy production Radioisotopes I-131 ◦ Diagnosing and treating thyroid disorders Co-60 ◦ Treating cancer C-14 ◦ Dating once-living organisms ◦ Compare to C-12 U-238 ◦ Dating geologic formations ◦ Compare to Pb-206 Half Life Amount of time for half of a sample to decay into a new element Fraction 1 Re maining 2 t # HalfLives T Mass Left Original Mass = 1 2 t T t T Example How many half lives does it take for a sample of C-14 to be 11430 yrs old? t 11430 y 2 T 5715 y More Practice How much 226Ra will be left in a sample that is 4797 years old, if it initially contained 408g? x 1 408g 2 4797 y 1599 y 3 1 1 2 8 51g And One More…. What is the half life of a sample that started with 144g and has only 9g left after 28days? 9g 1 144 g 2 28d x 1 1 16 2 4 28d 4 x 7d Main Menu Unit 4 - Periodic Table Chapter 6 History Dmitri Mendeleev developed the first Periodic Table ◦ based on increasing atomic mass. ◦ elements with similar chemical properties next to each other ◦ out of order by atomic mass if chemical properties lined up ◦ left open spaces for elements not yet discovered ◦ predicted properties of elements not yet discovered Periodic Table The modern periodic table is arranged in order of increasing atomic number. Arrangement Rows are called Periods Columns are called Groups Numbered 1-18 ◦ ◦ ◦ ◦ ◦ ◦ Group 1 - Alkali Metals Group 2 - Alkaline earth metals Group 17 – Halogens Group 18 - Inert or Noble gases Groups 3-11 – Transition Metals Bottom 2 rows – Inner Transition Phases at STP Most elements are solids at STP Hg and Br are liquids at STP H, N, O, F, Cl and Noble Gases are all gases at STP Valence Electrons Electrons in outermost occupied energy level Valence Electrons are responsible for most chemical properties ◦ Elements in the same group have similar properties because they have the same number of valence electrons Metals Good conductors of heat and electrical current High luster or sheen Many are ductile, meaning they can be drawn into wires Most are malleable, meaning they can be hammered into thin sheets Metals Metallic Character increases as you move towards the lower left Most Metallic Element is Francium, Fr Nonmetals Most are gases at room temperature, some are solids, and one is liquid Most are poor conductors Most solids are brittle Nonmetals Non-Metallic Character increases as you move towards upper right Most nonmetallic element is Fluorine, F Metalloids B, Si, Ge, As, Sb, Te Have properties of both metals and nonmetals, based on conditions Exceptions: ◦ Al and Po are metals ◦ At is a nonmetal Group Characteristics Alkali Metals (Group 1) ◦ H, Li, Na, K, Rb, Cs, Fr ◦ All have 1 valence electron, tend to form +1 ions ◦ Most reactive metals ◦ Not found in nature by themselves, always combined with someone else ◦ Have properties of metals but are softer and less dense Group Characteristics (cont) Alkaline Earth Metals (Group 2) ◦ Be, Mg, Ca, Sr, Ba, Ra ◦ All have 2 valence electrons, tend to form +2 ions ◦ Harder and more dense than alkali metals, but also have higher melting and boiling points ◦ Highly reactive, but not as much as alkali metals ◦ Not found by themselves in nature Group Characteristics (cont) Halogens (Group 17) ◦ F, Cl, Br, I, At ◦ All have 7 valence electrons, tend to form -1 ions ◦ Strongly non-metallic ◦ Most active nonmetals ◦ Have low melting and boiling points ◦ Combine readily with metals to form salts Group Characteristics (cont) Noble Gases (Group 18) ◦ He, Ne, Ar, Kr, Xe, Rn ◦ Colorless gases that are extremely nonreactive ◦ Full valence shell, non-reactive ◦ All are found in small amounts in our atmosphere Group Characteristics (cont) Transition Metals (Groups 3-11) ◦ Most are excellent heat and electrical conductors ◦ Most have high melting points and are hard, except Hg ◦ Less active than group 1 and 2 metals ◦ Many combine with Oxygen to form oxides (Chemical property) ◦ Many have more than one oxidation number ◦ Form compounds that are colorful Diatomics Eight elements are diatomic molecules when alone in nature (exist as two atoms bonded together) H2, N2, O2, F2, Cl2, Br2, I2, At2 Diatomics Hydrogen and the Magic 7 Trends Atomic number ◦ increases across a period. ◦ increases down a group Atomic mass ◦ generally increases across a period. ◦ increases down a group. Properties Atomic Radius – size of the atom Ion ◦ Atom, or group of atoms, that has gained or lost electrons Cation – positive ion Anion – negative ion Ionic Radius – size of an ion Ions Positive element Ions ◦ have lost electrons ◦ radius becomes smaller ◦ Metals tend to lose electrons Negative element Ions ◦ have gained electrons ◦ radius becomes larger ◦ Nonmetals tend to gain electrons Trends Atomic Radius ◦ decreases across a period ◦ increases down a group Ionic Radius ◦ decreases across a period for positive ions ◦ decreases across a period for negative ions ◦ increases down a group Properties Ionization Energy ◦ Amount of energy required to remove an electron from an atom Electronegativity ◦ Ability of an atom to attract an electron from another atom when in a compound. Noble gases are usually omitted since they don’t form compounds Trends Ionization Energy ◦ tends to increase across a period ◦ tends to decrease down a group Electronegativity ◦ tends to increase across a period ◦ tends to decrease down a group Trends Summary Property Atomic Number Atomic Mass Atomic Radius Ionic Radius Ionization Energy Electronegativity Period (LR) Group (TB) Main Menu Unit 5 - Bonding Chapters 7-8 Octet Rule Atoms tend to lose or gain electrons to achieve a full valence shell (8) ◦ Exception: First Energy Level is full with 2 electrons Electron Dot Structures Diagrams that show valence electrons, usually as dots ◦ AKA Lewis Electron Dot Diagrams Rules ◦ ◦ ◦ ◦ Start on any side First two get paired together Next three are separated Fill in as needed O Ions Atoms that have gained or lost electrons, and now have a charge Must show charge + Na - F -2 O Compounds Two Main Types of Compounds ◦ Ionic ◦ Molecular (Covalent) Based on type of bonding involved Bonding Bond ◦ Shared or exchanged electrons that hold two atoms together Three Main Types ◦ Covalent ◦ Ionic ◦ Metallic Covalent Bonds Electrons are shared between two atoms to hold them together ◦ Each atom will try to achieve a full valence shell ◦ 2 nonmetals Two types of covalent bonds ◦ Non-Polar Covalent – Shared equally ◦ Polar Covalent – Shared unequally Covalent Bonding HH O O N N OH H More Examples H Cl H H C H H H N H H O C O Determining Bond Type Based on electronegativity DIFFERENCE between two bonding atoms Nonpolar Covalent Bond ◦ 2 same Nonmetals (no difference in electronegativity) Polar Covalent Bond ◦ 2 different Nonmetals (small difference in electronegativity) Determining Bond Polarity The larger the difference in electronegativity, the more polar the bond. Bonding Ionic Bond ◦ Electrons are transferred from one atom to another (one gives, one takes) ◦ Metal and nonmetal, NaCl Large electronegativity difference Properties Ionic Compounds ◦ Most ionic compounds are hard, crystalline solids at room temperature ◦ High melting points ◦ Mostly soluble in water ◦ Can conduct an electric current when melted or dissolved in water(aq). Properties Covalent Compounds ◦ Most molecular compounds tend to have relatively lower melting and boiling points than ionic compounds. Ionic Compounds Ionic compounds are electrically neutral, even though they are composed of charged ions ◦ Total positive charge equals total negative charge Determining Formulas Must be electrically neutral ◦ Total positive charge must equal total negative charge Use oxidation numbers from Periodic Table ◦ Group 1 +1 ◦ Group 13 +3 ◦ Group 16 -2 Group 2 +2 Group 15 -3 Group 17 -1 Determining Formulas Determine number of each ion to balance out charge ◦ Use as subscript for element symbol ◦ Ex: CaCl2, Na3PO4, Mg(NO3)2 ◦ Write Positive Ion First Formula must be smallest wholenumber ratio Short-cut (criss-cross method) Magnesium and Phosphate Mg+2 PO4-3 Mg3(PO4)2 Short-cut (criss-cross method) Magnesium and Carbonate Mg+2 CO3-2 Mg2(CO3)2 MgCO3 Must Simplify Polyatomics Compounds with polyatomic ions contain BOTH ionic and covalent bonds ◦ Example: NaNO3 Allotropes Two or more different molecular forms of the same element in the same physical state ◦ Different properties because they have different molecular structures ◦ O2 vs O3 ◦ Diamond and Graphite (carbon) Metallic Bonding Bonding within metallic samples is due to highly mobile valence electrons ◦ Free flowing valence electrons ◦ “Sea of Electrons” Network Solids All atoms in a network solid are covalently bonded together Network solids have very high melting and boiling points, since melting requires the breaking of many bonds throughout the compound. Some of the strongest materials known to man are network solids. Bond Energy When two atoms form a bond, energy is released ◦ Example: Cl + Cl Cl2 + energy Energy needs to be added to break a bond ◦ Example: Cl2 + energy Cl + Cl Structural Formulas Shared electrons are written as a line, unshared electrons are not written ◦ Each line represents 2 electrons O H H H Cl H O O N H H Molecular Polarity Polar Molecule ◦ one end of a molecule is slightly negative(δ-) and the other end is slightly positive(δ+). ◦ Asymmetrical charge distribution Nonpolar Molecule ◦ Can not be separated into different ends ◦ Symmetrical charge distribution Polar Molecule H2O ◦ Polar Covalent Bond ◦ Electrons shared Unequally δ- O Hδ Hδ + + More Examples HCl δ+ H Cl δ- δ- NH3 δ+ H H N H δ+ H Hδ + Another Example CH4 δ+ H δ+ H C H δ δ- + Nonpolar Molecule H δ+ Polarity Ionic Compounds are Ionic Nonpolar Covalent Bonds always indicate Nonpolar Molecules Polar Covalent Bonds ◦ Determine Symmetry “Like Dissolves Like” Polar and Ionic substances will dissolve in other Polar Substances Nonpolar substance will dissolve in other nonpolar substances Intermolecular Forces Intermolecular Forces of Attraction ◦ attraction between two molecules or ions that hold them together (not a bond) ◦ Determines melting and boiling points of compounds Stronger intermolecular forces, higher melting and boiling points Hydrogen Bonding Hydrogen bonded to N, O, or F, is attracted to the N, O, or F of another molecule. Not actual bond, just attraction Hydrogen “Bond” H F H F Main Menu Unit 6 – Naming and Moles Chapter 9-10 Naming Ions Positive Ions, cations, simply retain their name. ◦ Na+ Sodium Ion ◦ Mg2+ Magnesium Ion Negative Ions, anions, change ending of element to –ide ◦ Cl- Chloride Ion ◦ Br- Bromide Ion Naming Systems Binary Covalent System (Prefixes) ◦ 2 nonmetals/metalloids Stock System (Roman Numerals) ◦ Metal has more than one positive oxidation number Ionic System ◦ Everything else (for now) Naming Ionic Compounds Name positive ion first, then negative ion. ◦ NaCl Sodium chloride ◦ Mg(OH)2 Magnesium hydroxide Stock System We must indicate which charge the metal ion has, using Roman Numerals ◦ ◦ ◦ ◦ CuCl CuCl2 Cu+1 Cu+2 Copper(I)Chloride Copper(II)Chloride Roman Numeral is the charge of the metal ion Roman Numerals Cation Charge +1 +2 +3 +4 Roman Numeral I II III IV +5 +6 +7 +8 V VI VII VIII Binary Covalent Compounds Use when compound is 2 nonmetals ◦ Including metalloids Use a prefix system to indicate the number of atoms for each element Second element ends in –ide N2Cl3 ◦ Dinitrogen Trichloride Prefixes Number of atoms 1 2 Prefix monodi- 3 4 5 6 tritetrapentahexa- 7 8 heptaocta- Exceptions When there is only one atom of the first element, do not use mono- prefix. ◦ CO2 Carbon dioxide When an element starts with a vowel, drop any o or a at the end of a prefix ◦ CO Carbon Monoxide ◦ P2O5 Diphosphorus Pentoxide Mole 6.02 x 1023 ◦ Avogadro’s Number Number of representative particles in a mole 1 mol He = 6.02 x 1023 atoms 1 mol H2 = 6.02 x 1023 molecules Gram Formula Mass Mass of the formula in g/mol Simply add the atomic masses of each element in the formula together H2O = 1 + 1 + 16 = 18 g/mol ◦ Also known as gram atomic mass, gram molecular mass, molar mass Practice KNO3 = 39 + 14 + 16(3) = 101 g/mol C6H14 = 12(6) + 1(14) = 86 g/mol CuSO4 = 63.5 + 32 + 16(4) = 159.5 g/mol Mole - Mass Conversion Example: 96 g of Oxygen gas = ? mol 96 g # mol 3mol 32 g / mol Molar Volume At STP, 1 mol of any gas occupies 22.4L of space Examples: 2 mol of He occupies how much space at STP? 44.8L 11.2L will have how much H2 gas at STP? 0.5 mol Mole Road Map Percent Composition Part *100% Whole Another Example NH3 N H 14 *100 17 3 *100 17 82.4% N 17.6% H Hydrates Compounds that have a specific number of water molecules attached CuSO4·5H2O ◦ Dot means plus (+) ◦ gfm = 159.5 + 5(18) = 249.5g/mol Empirical Formula Simplest Whole-Number ratio of atoms in a compound ◦ Molecular Formula is a multiple of the Empirical Formula Examples CO2 P4O10 C6H12O6 P2O5 CH2O Empirical Formula A molecular formula has an empirical formula of CH2 and a molecular mass of 28 g/mol. C 2H 4 A molecular formula has an empirical formula of CH2 and a molecular mass of 42 g/mol. C 3H 6 Main Menu