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by: Regan Dougherty

CH104FinalStudyGuide.pdf CH 104

Regan Dougherty
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CH 104 final exam study guide. Includes chapters 1-10 with emphasis on 8.7 - 9.
Introductory Chemistry
Stephen Woski
Study Guide
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This 15 page Study Guide was uploaded by Regan Dougherty on Thursday December 3, 2015. The Study Guide belongs to CH 104 at University of Alabama - Tuscaloosa taught by Stephen Woski in Summer 2015. Since its upload, it has received 156 views. For similar materials see Introductory Chemistry in Chemistry at University of Alabama - Tuscaloosa.

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Date Created: 12/03/15
Exam Date: 12/8/15 CH 104 Final Exam Study Guide 18 questions on chapters 8 and 9. 4 questions for every other chapter. - Chapter 1: Matter and Measurement • Matter - anything that has mass and takes up volume - 3 States: • Solid - definite shape and volume • Liquid - definite volume; takes the shape of its container • Gas - no definite shape or volume • Pure substance - contains a single component and has a constant composition - Element - pure substance that cannot be broken down by a chemical reaction - Compound - pure substance formed by chemically combining two or more elements; can be broken down by a chemical reaction • Mixture - composed of more than one substance • Temperature Conversions - Celsius to Fahrenheit • F = 1.8(C) + 32 - Celsius to Kelvin • K = C + 273 • Density = mass / volume - Chapter 2: Atoms and the Periodic Table • Metals - good conductors of heat and electricity // found on the left side of the periodic table • Nonmetals - generally poor conductors of heat and electricity // found on the right side of the periodic table • Metalloids - have properties intermediate between metals and nonmetals • Structure of the Atom 1 Exam Date: 12/8/15 - Proton - positive charge - Electron - negative charge - Neutron - no charge - Nucleus - dense core that contains protons and neutrons - Electron cloud - composed of electrons that move almost rapidly in an empty space surrounding the nucleus - Isotopes - atoms of the same element having different number of neutrons - Every atom of the same element has the same number of protons in its nucleus. - Periodic Table • A period is a row and a group is a column. • Notable Groups: - Alkali metals (group 1A) - Alkaline earth metals (2A) - Halogens (7A) - very reactive - Noble gases (8A) - stable - Electronic Structure • Orbital - a region of space where the probability of finding an electron is high - holds 2 electrons • The electrons that surround a nucleus are confined to regions called the principle energy levels or shells. - Electrons closer to the nucleus are lower in energy. - Shell 1 holds 2 electrons. - Shell 2: 8 electons - Shell 3: 18 electrons - Shell 4: 32 electrons - The s subshell contains one s orbital. 2 Exam Date: 12/8/15 - The p subshell has 3 p orbitals. - The d subshell has 5 d orbitals. - The f subshell has 7 f orbitals. • Valence electrons - electrons in the outermost shell - Elements in the same group have the same number of valence electrons. This causes them to have similar properties. • Periodic Trends - The size of atoms increases down a column. - The size of atoms decreases across a row. • Ionization energy - the energy needed to remove an electron from a neutral atom - Ionization energy decreases down a column of the periodic table. - Ionization energy increases across a row. - Chapter 10: Nuclear Chemistry • Nuclear reaction - a reaction that involves subatomic particles of the nucleus • Types of Radiation - Alpha particle - 2 protons and 2 neutrons; mass of 4; charge of +2 • 42He low energy • - Beta particle - electron; mass of 0; charge of -1 0 • -1 • better penetrator than alpha - Positron - antiparticle of a beta particle; mass of 0; charge of +1 • noted as p11 - 11p —> 10n + 01e - Gamma ray - mass and charge of 0 3 Exam Date: 12/8/15 • high energy light • highest energy • Radioactive decay - the process by which an unstable radioactive nucleus emits radiation, forming a nucleus of a new composition • Half-life - the time it takes for half of a radioisotope sample to decay • RAD (radiation absorbed dose) - amount of radiation absorbed by one gram of a substance REM (radiation equivalent for man) - amount of radiation that also factors in its • energy and potential to damage tissue • Nuclear fission - the splitting apart of a heavy nucleus into lighter nuclei and neutrons • Nuclear fusion - he joining together of two nuclei to form a larger nucleus - Chapter 3: Ionic Compounds • Ionic bonds result from the transfer of electrons from one element to another. - occur between a metal and a nonmetal • Ionic compounds consist of oppositely charged ions that have a strong attraction for each other. • cations - positive ions • anions - negative ions • Octet rule - atoms like to have a noble gas configuration (usually 8 valence electrons) • The charge of transition metals is normally indicated by Roman numerals after the name. - Another way to indicate charge: ion with the lower charge end in -ous and the ion with the high charge end in -ic - copper(I) - cuprous - copper(II) - cupric - tin(II) - stannous 4 Exam Date: 12/8/15 - tin(IV) - stannic - iron(II) - ferrous - iron(III) - ferric • Polyatomic ion* - cation or anion that contains more than one atom held together by covalent bonds - Chapter 4: Covalent Compounds • Covalent bond - a two-electron bond in which the bonding atoms share valence electrons - occurs between two nonmetals or a nonmetal and a metalloid • Molecule - a group of atoms held together by covalent bonds • Molecular Shape - Linear - occurs when an atom is surrounded by two groups • bond angles = 180° - Trigonal planar - atom surrounded by three groups • bond angles = 120° - Tetrahedral - atom surrounded by four groups • bond angles = 109.5° - Lone pairs of electrons take up space just like bonds do. • Trigonal pyramidal - 3 atoms and one lone pair • Bent - 2 atoms and 2 lone pairs • Not all atoms in covalent bonds share electrons equally. - Electronegativity - a measure of an atom’s attraction for electrons in a bond • increases across a row • decreases down a column - Nonpolar bond - a bond where electrons are equally shared - Polar / polar-covalent bonds result from bonding between atoms with different electronegativities. 5 Exam Date: 12/8/15 - Dipole - partial separation of charge in a bond or molecule - Chapter 5: Chemical Reactions • Chemical reactions involve breaking bonds in the starting materials (reactants) and forming bonds in the products. • Law of conservation of matter - atoms cannot be created or destroyed in a chemical reaction • Oxidation - the loss of electrons Reduction - the gain of electrons • • A compound that gains electrons (is reduced) while causing another compound to be oxidized is called an oxidizing agent. • A compound that loses electrons (is oxidized) while causing another compound to be reduced is called a reducing agent. • Mole - 6.02 x 10 23 - Molar mass - the mass of one mole of any substance (same value as atomic weight) • Percent Yield - Theoretical yield - the amount of product expected from a given amount of reactant based on the coefficients in the balanced chemical equation - Actual yield - the amount of product isolated from a reaction - percent yield = actual / theoretical x 100% • Limiting reactant - the reactant that is completely used up in a reaction; controls how much of a product is produced - To determine limiting reactant: divide the number of moles you have of a substance by the coefficient (do this for all of the reactants). The smaller number is the limiting reactant. - Chapter 6: Energy Changes, Reaction Rates, and Equilibrium • Energy - the capacity to do work - Potential energy - stored energy • Chemical bonds store potential energy. 6 Exam Date: 12/8/15 - Kinetic energy - energy of motion • Law of conservation of energy - The total energy of the universe does not change. Energy cannot be created or destroyed. • A compound with lower potential energy is more stable than a compound with higher potential energy. • Calories and joules are units of energy. • Energy Changes in Reactions - Breaking a bond requires energy (endothermic). - Forming a bond releases energy (exothermic). - Heat of reaction (ΔH) - the energy absorbed or released in any reaction - ΔH = energy of products - energy of reactants • Endothermic - energy is absorbed (ΔH is +) - Products are higher in energy than the reactants. • Exothermic - energy is released (ΔH is -) - Products are lower in energy than the reactants. - Bond dissociation energy - the energy needed to to break a covalent bond - In order for molecules to react, they must collide. - Energy of activation (E )a- minimum amount of energy needed for a reaction to occur • Reaction Rates - Increasing the concentration of reactants increases the number of collisions, so the reaction rate increases. - Increasing the temperature increases kinetic energy, so the reaction rate increases. - Catalyst - a substance that increases the rate of a reaction but is recovered unchanged at the end of the reaction • lowers the E a • Equilibrium 7 Exam Date: 12/8/15 - A reversible reaction can occur in either direction. - When the rate of the forward reaction equals the rate of the reverse reaction, the net concentrations of all species do not change and the system is at equilibrium. - Equilibrium constant (K) = [products] / [reactants] • [ ] signifies concentration • The coefficients in the equation are used as exponents. • K = 1 means the reaction is at equilibrium. • Le Chatelier’s Principle - If a chemical system is disturbed or stressed, the system will react in the direction that counteracts the disturbance or relieves the stress. - Chapter 7: Gases, Liquids, and Solids • See Chapter 1 States of Matter for brief description of states • Gases and Pressure - The size of gas particles is negligible compared to the space between the particles. Gas particles do not exert attractive forces on one another. - Pressure - the force exerted per unit of area when gas particles hit a surface - Units of Pressure 1 atm = 760 torr = 760 mm Hg • - Combined gas law - (P V ) / T = (P V ) / T 1 1 1 2 2 2 Avogadro’s law - When pressure and temperature are constant, the volume of a • gas is proportional to the number of moles present. • Ideal gas law - PV = nRT - Temperature must be in K. - R is the universal gas constant. • Dalton’s law - The total pressure of a gas mixture is the sum of the partial pressures of its component gases. - Partial pressure - the pressure exerted by one component of a mixture of gases • Intermolecular forces - attractive forces that exist between molecules of liquids and solids 8 Exam Date: 12/8/15 - London dispersion forces (van der Waals interactions) - very weak interactions due to momentary changes in electron density in a molecule • occur in all covalent compounds - Dipole-dipole interactions - attractive forces between permanent dipoles of two polar molecules - Hydrogen bonds - occurs when a hydrogen atom bonded to an O, N, or F is elecrostatically attracted to an O, N, or F of another molecule • strongest type of intermolecular force Boiling point - the temperature at which a liquid is converted to the gas phase • • Melting point - the temperature at which a solid is converted to the liquid phase - The stronger the intermolecular forces, the higher the boiling/melting point. • Specific heat - the amount of energy it takes to raise the temperature of 1 g of a substance by 1° C • Calculating Heat to Warm/Cool a Single-Phase Sample - Heat = specific heat x mass x Δtemperature • Energy and Phase Changes - Melting and vaporization are endothermic processes. • Heat of fusion (ΔH fusion) - the amount of energy needed to melt one gram of a substance - heat required to convert solid to liquid = ΔH fusion x mass • Heat of vaporization (ΔH vaporization) - the amount of energy needed to vaporize one gram of a substance - heat required to convert gas to liquid = ΔH vaporization x mass - Freezing and condensation are exothermic processes. - Chapter 8: Solutions • Mixtures can be homogenous or heterogenous. - Homogenous - uniform composition throughout - Heterogenous - does not have a uniform composition throughout 9 Exam Date: 12/8/15 • Solution - homogenous mixture that contains small particles - Solute - substance present in lesser amount in a solution - Solvent - substance present in greater amount in a solution • Electrolyte - a substance that conducts an electric current in water - Strong electrolytes dissociate completely in water; weak electrolytes do not. - Non-electrolyte - a substance that does not conduct an electric current in water • “Like dissolves like.” (Nonpolar compounds are soluble in nonpolar solvents, etc.) - A compound is soluble in water if it contains one of the following cations: Group 1A (Li+, Na+, K+, Rb+, Cs+) • • Ammonium (NH +) 4 - A compound is soluble in water if it contains one of the following anions: • Halide: Cl-, Br-, I- (except for salts with Ag+, Hg 22+, and Pb )+ • Nitrate (NO -3 • Acetate (CH CO3-) 2 • Sulfate (SO )4(except for salts with Ba , Hg2+ 22+, andPb )2+ • Concentration - the amount of solute dissolved in a given amount of solution - Weight / volume percent (grams of solute / mL of solution) - Volume/volume percent (mL of solute / mL of solution) - Parts per million (ppm) - Molarity (moles of solute / L of solution) - The following material will be emphasized on the exam (Ch. 8.7 - Ch. 9). • Dilution - the addition of solvent to decrease the concentration of the solution - C1V 1 C V2 2 • Colligative Properties - properties of a solution that depend on the concentration of the solute but not its identity 10 Exam Date: 12/8/15 - Boiling Point Elevation - the increase in the boiling point of a liquid solution due to the presence of a dissolved nonvolatile solution • Volatile- readily vaporized • Nonvolatile - not readily vaporized • One mole of any nonvolatile substance raises the boiling point of one kilogram of water by 0.51° C. - Freezing Point Depression - he decrease in the melting point of a liquid solution due to the presence of a nonvolatile solute • One mole of any nonvolatile solute lowers the freezing point of one kilogram of water by 1.86° C. • Osmosis and Dialysis - Osmosis - the passage of a solvent across a semi-permeable membrane from a solution of low solute concentration to a solution of higher solute concentration - Osmotic pressure - the pressure that prevents the flow of additional solvent into a solution on one side of a semi-permeable membrane The greater the number of dissolved particles, the greater the osmotic • pressure. • Isotonic - describes two solutions with the same osmotic pressure - Solvent does not move. • Hypotonic - describes a solution that has a lower osmotic pressure than the surrounding solution - Solvent moves in. • Hypertonic - describes a solution that has a higher osmotic pressure than the surrounding solution - Solvent moves out. - Dialysis - a process that involves the selective passage of substances across a semi-permeable membrane - Chapter 9: Acids and Bases • Introduction to Acids and Bases 11 Exam Date: 12/8/15 - ACID - contains a hydrogen atom and dissolves in water to form a hydrogen ion (H+) • A Bronsted-Lowry acid is a proton donor. - must contain a hydrogen atom • monoprotic (contains one acidic proton), diprotic (two), or triprotic (three) acids - BASE - a base contains hydroxide and dissolves in water to form OH- • A Bronsted-Lowry base is a proton acceptor. - Because a proton has no electrons, a base must contain one lone pair of electrons that can be donated to form a new bond. - Acids to know: • Strong: - HCl (hydrochloric acid) - HNO (3itric acid) - H 2O (4ulfuric acid) • Weak: - H 2O (3arbonic acid) - H 3O (4hosphoric acid) - CH C3OH (acetic acid) - Bases to know: • Strong: - metal hydroxides (ex. NaOH) • Weak: - NH 3 • The Reaction of a Bronsted-Lowry Acid with a Bronsted Lowry Base - A proton is transferred from the acid to the base. • conjugate acid / base pair 12 Exam Date: 12/8/15 - In an acid-base reaction, one bond is broken and one bond is formed. - The net charge must be the same on both sides of the equation. - A compound that contains both a hydrogen and a lone pair of electrons can be either an acid or a base, depending on the reaction. This compound is said to be amphoteric. • Acid and Base Strength - Dissociation - the splitting apart of a covalent molecule or an ionic compound into individual ions • When a strong acid dissolves in water, 100% of the acid dissociates into ions. - Strong acids are more likely to donate protons than weak acids. - Strong acids form weak conjugate bases. • When a weak acid dissolves in water, only a small fraction of the acid dissociates into ions. • The same dissociation trends apply to strong and weak bases. - Strong bases form weak conjugate acids. - The position of equilibrium depends on the strength of the acids and bases. • When the stronger acid and base are the reactants, the reaction readily occurs and the reaction favors the products. • Equilibrium and Dissociation Constants - acid dissociation constant (K ) a K[H O]2= [H O+]3A: -] / [HA] • K a concentration of products / concentration of reactants • The stronger the acid, the larger the value of K . a • Equilibrium favors the formation of the acid with the smaller K (weaker acid). • The Dissociation of Water • K w [H O3][HO-] - neutral: [H 3+] = [HO-] - acidic: [H 3+] > [HO-] 13 Exam Date: 12/8/15 - basic: [H O+] < [HO-] 3 • The pH Scale - pH = -log[H O3] -pH - [H 3+] = 10 - Acid: pH < 7 - Neutral: pH = 7 - Base: pH > 7 - The lower the pH the higher the concentration of H O+. 3 - Since pH is measured on a logarithmic scale, a small difference in pH translates to a large change in H O+3concentration. - Physiological pH = 7.4 - Stomach pH = 1.7 - Small intestine pH = 8.5 - Blood pH = 7.4 • Common Acid-Base Reactions - The reaction of an acid with an OH- base is an example of a neutralization reaction (produces salt and water as products). • The acid donates a proton to the OH- base to form H O. 2 • The anion from the acid combines with the cation from the base to form a salt. • Net ionic equation: H + OH —> H O 2 - Net ionic equation - contains only the species involved in a reaction. • The Acidity and Basicity of Salt Solutions - A salt can form an acidic, neutral, or basic solution depending on whether its cation and anion are derived from a strong or weak acid and base. • The cation comes from the base and the anion comes from the acid. • strong base + weak acid = basic weak base + strong acid = acidic • 14 Exam Date: 12/8/15 Titration - a technique for determining an unknown molarity of an acid by adding a • base of known molarity to a known volume of acid - When the number of moles of bases added equals the number of moles of acid in the flask, the acid is neutralized. • Buffer - a solution whose pH changes very little when acid or base is added - consists of a weak acid and its conjugate base - to calculate the pH of a buffer solution: [H 3+] = K xa( [acid] / [conjugate base] ) *Check out Quizlet for polyatomic ion flash cards. My username is rmdougherty1 - use of the flash cards is completely free! If you have any questions feel for to contact me at or or though Blackboard messaging. Good luck! 15


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