SAT Review Flashcard

1. Mass

  • amount of matter in a sample
    • matter = anything that occupies space & has mass
  • grams, milligrams, kilograms
  • more mass means more atoms/molecules

1. Volume

  • how much space sometimg takes up
  • usually liters or milliliters
  • 1 cm3 = 1 mL
  • measure liquid volume with measuring flask
  • measure solid volume with displacement
  • gas volume = volume of container (displacement with container if necessary)

1. Density

  • ratio of mass to volume of an object
  • each substance has a specific density
  • d = m/v
  • density of liquids and solids as a fixed temp is constant
  • density of gases changes

1. Presh

  • force that gas in a closed container exerts on the container walls
  • solid/liquid is there, gas exerts presh on walls of environment and everything in it including the solid/liquid
  • use manometer or barometer to measure (both use Hg)
  • 760 torr = 760 mmHg = 1 atm

1. Energy

  • the ability to do work or transfer heat
  • heat, light, kinetic, chemical bond energy (different forms)
  • usually kinetic energy of molecules in chem–greater KE means higher temp and faster movement
  • breaking bonds takes energy, forming them releases energy
  • heat is the transfer of KE from one thing to another
  • 1 cal = 4.186 J; 1000 J = 1 kJ
  • measure it with a calorimeter

1. Temperature and Heat

  • heat is energy flow from high temp thing to low temp thing
  • temp measures average KE of molecules in a sample
  • if something increases in temp, that’s a heat content increase
  • substances have different specific heats
  • K = C + 273; 0 K = -273 C; 0 C = 273 K

1. Specific Heat

  • it’s easier to heat some things than others
  • heat capacity is how much energy something has to absorb for some of it to be raised 1 degree Celsius
  • specific heat is the heat capacity for 1 g of the substance
  • q = mst

2. Atom

  • smallest particle of an element
  • retains the chemical properties of the element

2. Parts of an Atom

  • proton, positive charge
  • neutron, no charge
  • nucleons, in the nucleus, protons and neutrons
  • electrons, negative charge
  • nucleons have mass; electrons have practically none
  • neutrons determine the isotope

2. Ion

  • atom with unequal charges inside and outside nucleus
  • cation, positively charged
  • anion, negatively charged

2. Element

  • most fundamental unit of matter
  • can’t break it down without losing its identity

2. Table Organization

  • horizontal rows are periods
  • vertical columns are groups
  • all elements in the same group have the same number of electrons in outer shells and share properties

2. Atomic Number

  • number of protons in the nucleus of an atom
  • specific to the element

2. Mass Number

  • protons and neutrons have 1 amu mass
  • sum of an atom’s proton and neutrons is the mass number
  • isotopes have different numbers of neutrons in their nuclei
  • doesn’t appear on the table because it varies

2. Atomic Weight

  • average mass number, so it’s based on each isotope’s natural abundance
  • one for each element
  • listed on the table

3. Molecule

  • units of two or more atoms
  • held together with chemical bonds

3. Diatomic Molecule

  • molecule made of just two atoms
  • some elements exist this way at stp: O2, I2, H2, N2, Cl2, F2, Br2
  • Clearly I Have NO Friends, Bro

3. Formula Weight

  • add the atomic weights of all the atoms in the molecule

3. Empirical Formula and Molecular Formula

  • empirical formula is the ratio of atoms within a molecule
  • the molecular formula divided by the gcf of the subscripts gives the empirical formula

3. Percent composition

  • refers to percent by mass
  • find it by calculating stuff 😛

3. Mole

  • a mole of something is 6.022 x 1023
  • the 6 number is called Avogadro’s number
  • a mole of atoms make up the atomic mass

3. Mass Composition to Empirical Formula

  • imagine 100g
  • convert percents to grams
  • convert grams to moles
  • find the ratio between the mole amounts

3. Stoichiometry Problems

  • how much product or reactant is produced or needed in a reaction?
  • always start with a balanced equation
  • remember limiting reactants
  • convert things to moles!

3. Entropy

  • S symbol
  • If deltaS is negative, reaction loses entropy
  • If deltaS is positive, reaction gains entropy
  • Universe tends towards disorder
  • Low energy states ar emore stable than high energy states
  • Higher S + Lower E = More Stable

3. Enthalpy

  • term refers to energy states of reactions or products
  • symbolized with H
  • decreasing H leads to stability
  • H decreases: exothermic; H increases: endothermic
    • endothermic reactions need energy input
  • energy of products/reactions refers to the energy in the bonds

3. Heat of Formation

  • amount of heat released or absorbed when a mole of a compound is formed
  • if Hf is negative, then exothermic, and vice versa
  • for all elements, the heat of formation is zero
  • heat of formation for entire reaction = sum of product Hfs minus sum of reactant Hfs
  • Hess’s Law

3. Hess’s Law

  • If a reaction happens in more than one step, the change in enthalpy for the whole thing is the sum of the changes in enthalpy for each step
  • Enthalpy is pathway-independent: state function

3. Spontaneous Reaction

  • happens without energy input
  • can happen for a positive H if S is low enough
  • can happen for a positive S if H is low enough
  • spontaneity is determined by Gibbs free energy

3. Gibbs Free Energy

  • ΔG =  ΔH – TΔS
  • T is in Kelvins
  • determines if a reaction will happen spontaneously or not
  • if ΔG < 0, then it’s spontaneous in that direction
  • if;;G ; 0, then it’s spontaneous in the opposite direction
  • if ΔG = 0, then there’s equilibrium

4. Quantum Mechanics

  • our current theory about how electrons and atoms work
  • a quantum is a small unit of energy, and all energy is quantized, so all energy exists in multiples of quanta

4. Orbitals

  • electrons exist in them
  • replaces Bohr model of orbiting
  • describes the likelihood that an electron will be found in a particular location–a probability function
  • orbitals have energy shells, shape, subshells
  • any orbital holds two electrons

4. Energy Shell

  • each orbital is in an energy shell
  • higher number energy shell
    • greater energy
    • farther from nucleus
  • each energy shell has a whole number

4. Orbital Shape and Subshell

  • 4 different shapes that make up subshells
  • s has 1 orbital
  • p has 3 orbitals
  • d has 5 orbitals
  • f has 7 orbitals

4. Heisenberg Principle

  • it’s impossible to know both the position and momentum of an electron at the same time

4. Louis De Broglie

  • matter has the properties of a wave
  • electrons behave in waves like electromagnetic radiation

4. Bohr Model

  • wrong
  • electrons orbit the nucleus in true orbits, like how planets orbit the sun

4. Aufbau Principle

  • each subshell is filled before electrons go in the next subshell
  • Cr and Cu are exceptions (promote a 4s to a 3d)
  • Hybrid orbitals are exceptions

4. Radioactivity

  • Unstable nucleus undergoes nuclear decay –> radioactive –> releases radioactivity
  • radioactive particles
    • alpha
    • beta
    • gamma
  • as a radioactive atom decays
    • becomes another isotope or
    • becomes another element

4. Geiger Counter

  • Detects and measures radioactive particles

4. Four Types of Radioactive Decay

  • alpha decay
  • beta decay
  • positron emission
  • gamma decay

4. Alpha Decay

  • emits alpha particle: 2 pros and 2 neus
  • decreases atomic # by 2 and atomic mass by 4
  • sometimes symbolized as 4/2He

4. Beta Decay

  • emits a beta particle: one electron
  • releases an electron from a neutron –> lose neutron and gain proton
  • atomic # increases by 1 and mass stays the same

4. Positron Emission

  • releases positron: antiparticle of an electron, same magnitude but opposite charge
  • proton becomes a neutron
  • decreases atomic # by 1, mass stays the same

4. Gamma Decay

  • emits gamma rays when the atom has too much energy
  • makes nucleus more stable but doesn’t do anything else

4. Half-life

  • rate of a substance’s radioactive decay
  • takes one half-life for half of the substance to decay

5. Horizontal and Vertical Similarities

  • same row (period) –> electrons in same energy shells
  • same column (group) –> similar chemical and physical properties

5. Family Traits

  • same column
  • noble gases
  • alkali metals
  • alkaline earth metals
  • halogens

5. Alkali Metals

  • group 1/1A
  • 1 valence electron
  • very reactive
  • snihy, grayish-white
  • melt easily
  • lower densities than other metals

5. Alkaline Earth Metals

  • group 2/2A
  • 2 valence electrons
  • less reactive than alkali metals but more reactive than other metals

5. The Active Metals

  • Alkali metals
  • Alkaline earth metals
  • they’re really reactive…duh

5. Halogens

  • group 17/7A
  • 7 valence electrons
  • very reactive
  • look different from each other

5. Metals

  • shiny and conduct heat well
  • malleable and ductile
  • all solid at room temp except mercury
  • lose electrons in bonds

5. Nonmetals

  • share or gain electrons in bonds
  • poor conductors of heat and electricity
  • some are solid, liquid, gas at room temp

5. Semimetals/Metalloids

  • some characteristics of metals and nonmentals
  • can either gain or lose electrons in a bond

5. Periodic trends to know

  • Ionization energy
  • Electronegativity
  • Atomic radius
  • Metallic character

5. Atomic Radius

  • increases towards cesium

5. Ionization Energy

  • energy required to remove an electron from an atom
  • increases towards fluorine
  • 1st, 2nd, etc.

5. Electronegativity

  • pull that a nucleus exerts on electrons of another atom
  • increases towards fluorine

5. Metallic character
5. Lattice Energy

  • energy required to completely separate a mole of a solid ionic compound into separate ions
  • higer lattice energy = stronger bond

5. Electrostatic Force

  • attraction betweena positive charge and a negative charge
  • strong

5. Transition Metals

  • middle of the periodic table
  • harder than metals
  • higher melting point than metals
  • compounds with transition metals are colorful

6. Ideal Gas Assumptions

  • molecules of an ideal gas don’t attract or repel each other
  • molecules of an ideal gas occupy zero volume

6. Kinetic Molecular Theory

  • kinetic energy of a gas molecule increases proportionally with temperature in K

6. Network Solid

  • atoms bonded by covalent bonds in a continuous network
  • one big macromolecule
  • no individual molecules
  • e.g. diamonds or quartz

6. Hydrate

  • Ionic substance where water molecules bond to the ions in a fixed ratio
  • dot H2O means it’s a hydrate
  • % comp of water by mass in a hydrate is called water of hydration

6. Four Types of Crystalline Solids

  • Ionic
  • Covalent network
  • Molecular
  • Metallic

6. Heat of Fusion

  • amount of heat it takes for a substance to go from solid to liquid

6. Heat of Vaporization

  • Amount of heat it takes for a substance to go from liquid to gas

6. Phase change and presh

  • more pressure means it’s harder to get solids to melt or liquids to vaporize
  • more presssure does the opposite for water–easier to melt

6. Triple Point

  • a particular temperature and pressure
  • substance can exist as a solid, liquid, or gas

6. Vapor Presh

  • created when liquids below their boiling points evaporate
  • all liquids in a closed system exert some
  • higher vapor presh means more volatility
  • affected by intermolecular forces, temperature, and molecular weight
  • boiling
    • all of the presh above a liquid is vapor presh
    • vapor presh = atmospheric presh

6. Energy and Phase Change

  • solid ———–> gas
  • low PE ——–> high PE, which universe dislikes
  • low entropy –> high entropy, which universe likes
  • melting/boiling is spontaneous when the temperature is above the melting/boiling point

7. Molarity

  • M
  • M = moles of solute/liters of solution

7. Molality

  • m
  • m = moles of solute/kilograms of solvent

7. Solubility Factors

  • Solids
    • more soluble at higher temperatures
  • Gases
    • less soluble at higher temperatures
    • more soluble at higher presh
    • think about carbonated drinks

7. Electrolytes

  • ions in solution
  • makes the solution about to conduct electricity
  • the solution is still neutral
  • in ionic solutions aka elecrolytic solutions

7. Solutes and Freezing/Boiling Points

  • ΔT = kmi
  • solute effects an increase in boiling point and decrease in freezing point
  • k changes from solvent to solvent
  • m = molality
  • i = the number of particles it dissolves into per mole
    • sucrose: i = 1
    • NaCl: i = 2

7. Precipitate

  • solid substance that settles out of a solution
  • often the result of a double replacement reaction
  • use solutbility rules

7. Solubility Rules

  1. all nitrates and perchlorates are soluble
  2. all alkali metal and ammonium compounds are soluble
  3. hydroxides are insoluble except alkali metals and BaOH
  4. silver, lead, and mercury salts are insoluble except with nitrates and perchlorates

8. Kinetics

  • study of the rates of reactions

8. Equilibrium

  • the point in a chemical reaction at which concentration of reactants and products ceases to change
  • forward and reverse reaction rates are equal

8. Activated Complex/Transition State

  • unstable
  • place where products form
  • need to reach activation energy before activated complex can form

8. Factors that affect likelihood of reactions

  • frequency of collision
    • concentration of reactants
    • surface area of reactants
    • temperature
  • energy of collision
    • temperature
    • nature of reactants
    • catalysts

8. Catalysts

  • increase rate of reaction by lowering activation energy
  • not consumed in a reaction
  • don’t change the equilibrium of a reaction

8. Equilibrium constant expression

  • reaction: aA + bB <—> cC + dD
  • Keq = ([C]c [D]d)/([A]a [B]b)
    • solvents not included
    • powers are coefficients from the equation
  • Keq = products/reactants
  • if > 1 then forward reaction favored
  • if < 1 then reverse reaction favored

8. Le Chatelier’s Principle

  • add more on one side –> equilibrium shifts to other side
  • Keq stays constant unless temperature changes
  • if stress is placed on a reaction at equilibrium, then the equilibrium will shift in a direction that relieves the stress

8. Haber Process

  • makes ammonia

8. Pressure Change Effect on Equilibrium

  • only applies if some of the things in the reaction are gases
  • reduce volume –> increase pressure –> equilibrium shifts to side with fewer moles of gas

8. Solubility Product Constant

  • Ksp
  • Equilibrium constant between insoluble ionic solid and dissolved ions
  • The values are usually really small
  • Solids aren’t included in normal equilibrium expressions

9. Autoionization

  • spontaneous dissociation of water tha tmakes H+ and OH- ions
  • reversible
  • an equilibrium exists between H+, OH-, and H2O
  • equilibrium expression for it is 10-14 M2 at 25 degrees C
    • same in every aqueous solution

9. pH

  • “p” means -log
  • if [H+] = 10-7 M, then pH = 7

9. Different Definitions for Acid and Base

  • Arrhenius
    • Acids make H+ in aq
    • Bases make OH- in aq
  • Lewis
    • acids accept electrons in aq
    • bases donate electrons in aq
  • Bronsted-Lowry
    • acids donate protons (H+)
    • bases accept protons
    • most common definition currently

9. Acid Dissociation Equations

  • HA(aq) –> H+(aq) + A(aq)
  • or
  • HA(aq) + H2O(l) –> H3O+(aq) + A(aq)

9. Base Reaction Equations

  • A(aq) + H+(aq) –> HA(aq)
  • or
  • A(aq) + H2O(l) –> HA(aq) + OH(aq)

9. Amphoteric

  • molecules or ions that can act as acids or bases depending on the solution
  • most molecules are either acids or bases

9. Strong Acids and Bases

  • completely dissociate
  • dissociation is 100% irreversible
  • one-way arrow
  • pH of 1.0 M strong acid is always 0
  • pH of 1.0 M strong base is always 14

9. List of Strong Acids

  1. HCl
  2. HBr
  3. HI
  4. HNO3
  5. H2SO4 (only the first H is strong)
  6. HClO4

9. List of Strong Bases

  1. LiOH
  2. NaOH
  3. KOH
  4. other group 1 hydroxides

9. Weak Acids and Bases

  • partially and reversibly dissociate
  • use the reversible arrow
  • figure out if something’s acidic or basic, then it’s weak if it isn’t on the strong list
  • remember that pH + pOH = 14 at 25 degrees C

9. Conjugate Pairs

  • molecules that are the same except one has an extra H+
  • the one with the extra H+ is the onjugate acid
  • the one with fewer H+ is the conjugate base
  • doesn’t necessarily mean that they’re an acid and a base
  • there’s an ionic equilibrium between the conjugates
  • sum of pKa and pKb of conjugates always = 14 at 25 C

9. Conjugate Rules

  1. Conjugate acid of a strong base is neutral
  2. Conjugate base of a strong acid is neutral
  3. Conjugate acid of a weak base is an acid
  4. Conjugate base of a weak acid is a base

9. Buffer

  • minimize a change in pH when an additional acid or base goes into a solution
  • a conjugate pair of a weak acid and a weak base so they don’t neutralize each other
  • use Henderson-Hasselbalch equation to calculate pH

10. Oxidation States

  • Sum of oxidation states in a compound is zero
  • Oxygen: -2
  • Alkali metals: +1
  • Alkaline Earth Metals: +2
  • Halogens: -1
  • Hydrogen: ±1

10. Oxidizing Agent

  • gets reduced
  • causes something else to get oxidized

10. Reducing Agent

  • gets oxidized
  • causes someone else to be reduced

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