Section 1
Preview this deck
mass number
Front
| 5 | 0 | ||
| 4 | 0 | ||
| 3 | 0 | ||
| 2 | 0 | ||
| 1 | 0 |
Active users
0
All-time users
0
Favorites
0
Last updated
6 years ago
Date created
Mar 1, 2020
Cards (137)
Section 1
(50 cards)
mass number
The sum of an atom's neutrons and protons
Energy Levels
s-subshell holds two electrons p-subshell holds six electrons d-subshell holds 10 electrons f-subshell holds 14 electrons
Ionic Bonds
An ionic solid is held together by the electrostatic attractions between ions that are next to one another in a lattice structure. Occurs between a metal and a nonmetal; electrons are not shared, they are given up by one atom and accepted by another. Substances with ionic bonds are usually solids at room temperature and have very high melting and boiling points. Ex. NaCl
Substitutional alloy
Forms between atoms of similar radii Ex. Atoms of zinc are substituted with copper atoms to create an alloy
Gases will most likely act as ideal under what conditions?
High temperature and low pressure
Network (Covalent) Bonds
In a network solid, atoms are held together in a lattice of covalent bonds. They are very hard and have very high melting and boiling points. Ex. The most commonly seen network solids are compounds of carbon (such as diamond or graphite) and silicon (SiO₂ quartz)
Periodic Trends
Across the periods - atomic radius decreases - ionization energy increases - electronegativity increases Down the periods - atomic radius increases - ionization energy decreases - electronegativity decreases
Factors Affecting Melting Points of Ionic Substances
1. Charge on ions - a greater charge leads to a greater bond energy Ex. MgO will have a higher melting point than NaCl 2. Size of ions - smaller ions will have greater attraction Ex. LiF will have a greater melting KBr
Moles
grams/molar mass
electron configuration
The complete description of the energy level and subshell that each electron on an element inhabits
Different types of bonds and their relative melting and boiling points (from highest to lowest)
1. Network Covalent Bonds 2. Ionic Bonds (based on Coulombic attraction) a. Greater Ion Charge b. Smaller Atom Size 3. Covalent Bonds (based on molecular polarity) a. Hydrogen Bonds b. Non-Hydrogen Bond Dipoles c. London Dispersion Forces (temporary dipoles) i. Large molecules are more polarizable because they have more electrons.
Atomic Radius Trends
Atomic radius decreases across a period Atomic radius increases down a group Cations are smaller than their atoms Ions are larger than their atoms
Incomplete Octets
Some atoms are stable with less than eight electrons in their outer shell. - Hydrogen only requires two electrons - Boron is considered to be stable with only six electrons, as in BF₃
Double Bond
Bond designation: One sigma and one pi Bond order: Two Bond length: Intermediate Bond energy: Intermediate
London Dispersion Forces
Forces that occur between all molecules. These very weak attractions occur because of the random motions of electrons on atoms within molecules. At a given moment, a nonpolar molecule might have more electrons on one side than on the other, given it an instantaneous polarity. Molecules with more electrons will experience will experience greater London dispersion forces, and therefore have generally higher melting and boiling points. London forces are even weaker than dipole-dipole forces, so substances that have only London dispersion forces melt and boil at extremely low temperatures and tend to be gases at room temperature.
Aufbau principle
States that when building up the electron configuration of an atom, electrons are placed in orbitals, subshells, and shells in order of increasing energy.
Hydrogen Bonds
Much stronger than dipole-dipole forces. Substances that have hydrogen bonds have higher melting and boiling points. Ex. water, H₂O and ammonia, NH₃ *Water is less dense as a solid than as a liquid because its hydrogen bonds force the molecules in ice to form a crystal structure, which keeps them apart than they are in the liquid form
Converting from moles to liters
I mole of gas = 22.4 L
Triple Bond
Bond designation: One sigma and two pi Bond order: Three Bond length: Shortest Bond energy: Greatest
ionization energy
The amount of energy necessary to remove an electron from an atom.
Hund's Rule
States that when an electron is added to a subshell, it will always occupy an empty orbital if no one is available. Electrons always occupy orbitals singly if possible and pair up only if no empty orbitals are available.
How intermolecular forces affect the phase of a substance
1. Substances with weak intermolecular forces (LD), tend to be gases at room temperature. 2. Substances with strong intermolecular forces (HB) tend to be liquids at room temperature. 3. Because ionic bonds are generally significantly stronger tan intermolecular forces in covalent molecules, ionic substances are usually solid at room temperatures. Ionic substances do not experience intermolecular forces.
Moles and Solutions
Moles = (molarity)(liters of solution)
Ionization energy
The energy required to remove an electron from an atom.
Avogadro's number
6.022×10²³ particles per one mole
Interstitial alloys
Metal atoms with two different radii combine Ex. In steel, much smaller carbon atoms occupy the interstices of the iron atoms
empirical formula
- Represents the simplest ratio of one element to another in a compound - Start by assuming a 100 g sample - Convert percentages to grams - Convert grams into moles - Divide each mole value by the lowest of the values - These values become the subscripts
molecular formula
- Determine the molar mass of the empirical formula - Divide that mass into the molar mass x = m/e x = molar mass/ empirical mass - Multiply all subscripts in the empirical formula by the value of x
Valence Shell Electron Pair Repulsion Model
Electrons repel each other, so when atoms come together to form a molecule, the molecule will assume the shape that keeps its different electron pairs as far apart as possible.
Electronegativity
Refers to how strongly the nucleus of an atom attracts the electrons of other atoms in a bond.
atomic number
The same as the number of protons in the nucleus of an element; it is also the same as the number of electrons surrounding the nucleus of an element when it is neutrally charged.
Formal Charge
If more than one valid Lewis structure exists for a molecule, formal charge can be used to determine the more likely structure.
Frequency and Wavelength
c = λv Inversely proportional c = speed of light in a vacuum (2.998×10⁸ m/s) λ = wavelength of the radiation v = frequency of the radiation
isotopes
Atoms of an element with different numbers of neutrons
Standard Temperature and Pressure (STP)
Pressure = 1 atm Temperature = 273 K
photoelectron spectra (PES)
A chart of the amount of ionization energy for all electrons ejected from a nucleus. The y-axis describes the relative number of electrons that are ejected from a given energy level. The x-axis shows the binding energy of those electrons.
Standard Enthalpy of Formation
The amount of heat lost or gained when one mole of a compound is formed from its constituent elements.
Expanded Octets
In molecules that have d subshells available, the central atom can have more than eight valence electrons, but never more than twelve. - PCl₅ - SF₄ - XeF₄
Heisenberg Uncertainty Principle
States that it is impossible to know both the momentum of an electron at a particular instant.
Enthalpy Change (Hess's Law for ΔH)
The enthalpy change for a reaction is equal to the sum of the enthalpy of formation of all the products minus the sum of the enthalpy of formation of all the reactants.
Energy and Electromagnetic Radiation
ΔE = hv = hc/λ ΔE = energy change h = Planck's constant, 6.626×10⁻³⁴ J∙s v = frequency of the radiation λ = wavelength of the radiation c = the speed of light, 3.00×10⁸ m/s
Covalent bonding
Bonding in which two atoms share electrons. Each atom counts the shared electrons as part of its valence shell to achieve complete outer shells. The first covalent bond formed between two atoms is called a sigma bond.
percent composition (mass percents)
The percent by mass of each element that makes up a compound. It is calculated by dividing the mass of each element or component in a compound by the total molar mass for the substance.
Resonance Forms
Single bonds
Bond designation: One sigma Bond order: One Bond length: Longest Bond energy: Least
Quantum Theory
Max Planck figured out that electromagnetic energy is quantized. That is, for a given frequency of radiation (or light), all possible energies are multiples of a certain unit of energy, called a quantum (mathematically, that's E = hv). So, energy changes do not occur smoothly but rather in small but specific steps.
Vapor Pressure
Arises from the fact that molecules inside a liquid are in constant motion. If those molecules hit the surface of the liquid with enough kinetic energy, they can escape the intermolecular forces holding them to the other molecules and transition them into the gas phase. This is not to be confused with a liquid boiling. In order for vaporization to occur, no outside energy needs to be added.
Coulomb's Law
The amount of energy that an electron has depends on its distance from nucleus of an atom. While on the exam, you will not be required to mathematically calculate the amount of energy a given electron has, you should be able to qualitatively apply Coulomb's Law. Essentially, the greater the charge of the nucleus, the more energy an electron will have.
Dipole-Dipole Forces
Forces that occur when the positive end of one polar molecule is attracted to the negative end of another polar molecule. Molecules with greater polarity have higher melting and boiling points. Dipole-dipole attractions, however, are relatively weak, and these substances melt and boil at very low temperatures.
Pauli Exclusion Principle
States that the two electrons which share an orbital cannot have the same spin. One electron must spin clockwise, and the other must spin counterclockwise.
Section 2
(50 cards)
Oxidizing agent
Reactant that is reduced (gains electrons)
Tetrahedral Geometry
- Central atom has four electron pairs - Zero lone pairs - sp³ hybridization - 109.5° bond angle - Ex. CH₄, NH₄⁺, ClO₄⁻, SO₄²⁻, PO₄³⁻
Activation Energy
The amount of energy required to reach the transition state, the highest point on the graph. At this point, all reactant bonds have been broken, but no product bonds have been formed, so this is the point in the reaction with the highest energy and lowest stability.
Reducing agent
Reactant that is oxidized (loses electrons)
Redox in a Galvanic Cell
Oxidation takes place at the anode Reduction takes place at the cathode
Gibbs Free Energy
The amount of energy in a system that is available to do useful work. -ΔG is spontaneous +ΔG is nonspontaneous When ΔG = 0 the reaction is at equilibrium
Bent Geometry
- Central atom with three electron pairs - One lone pair - sp² hybridization - 120° bond angle - Ex. SO₂
Solutes and Solvents
"Like dissolves like" A basic rule to remember which solutes will dissolve in which solvents. Polar or ionic solutes (such as salt) will dissolve in polar solvents (such as water). Nonpolar solutes (such as oils) are best dissolved in nonpolar solvents. When an ionic substance dissolves, it breaks up into ions in a process called dissociation. Free ions in a solution are called electrolytes because they can conduct electricity.
Graham's Law
The rate of diffusion of a gas molecule is inversely proportional to the square root of that molecule's mass.
If volume is constant
As pressure increases, temperature increases
Square Planar
- Central atom has 6 electron pairs - Two lone pairs - d²sp³ hybridization - Ex. XeF₄, ICl₄⁻
Catalysts
Speed up a reaction by providing the reactants with an alternate pathway that has a lower activation energy. A catalyst lowers the activation energy, but it has no effect on the energy of the reactants, the energy of the products, or the ΔH of the reaction.
Mole Fraction
Mole fraction gives the fraction of moles of a given substance (S) out of the total moles present in a sample. Mole Fraction (Xa) = moles of substance S / total number of moles in solution
Endothermic Reactions
If the products have weaker bonds than the reactants, then the products have higher enthalpy than the reactants and are less stable; in this case, energy is absorbed by the reaction, or the reaction is exothermic.
Trigonal Planar Geometry
- Central atom with three electron pairs - Zero lone pairs - sp² hybridization - 120° bond angles - Ex. BF₃, SO₃, NO₃⁻, CO₃²⁻
Density
D = m/v
Molarity (M)
Expresses the concentration of a solution in terms of volume. M = moles of solute / liters of solution
The Ideal Gas Equation
PV = nRT R = .08206 Latm/molK
Galvanic Cell (Voltaic Cell)
In a galvanic cell, a favored redox reaction is used to generate a flow of current. Two half-reactions take place in separate chambers, and the electrons that are released by the oxidation reaction pass through a wire to the chamber where they are consumed in the reduction reaction. That's how the current is created. If the concentration of the products in a voltaic cell increases, the voltage decreases. If the concentration of the reactants increases, the voltage increases.
Crystalline Solids
Solid that has its atoms arranged in an orderly way.
Reduction
Electron gain
Entropy
A measure of molecular randomness, or disorder
Cathode
Where reduction occurs and the solution is becoming less positively charged, the positive cations from the salt bridge solution flow into the half-cell.
Buffer Solution
A solution consisting of a weak acid plus its conjugate base or a weak base plus its conjugate acid. This solution resists changes to its pH
First Order Rate Laws
Rate = k[A] The rate law for a first order reaction uses natural logarithms. The use of natural logarithms in the rate law creates a linear graph comparing concentration and time. The slope of the line is given by -k and the y-intercept is given by ln[A]₀
Deviations From Ideal Behavior
At low temperature and/or high pressure, gases behave in a less-than-ideal manner. This is because the assumptions made in the kinetic molecular theory become invalid.
Half-life
Describes the amount of time it takes for half of a sample to react.
T-Shaped Geometry
- Central atom has 5 electron pairs - Two lone pairs - dsp³ hybridization - Ex. ClF₃, ICl₃
Specific Heat Capacity
The amount of heat required to raise the temperature of one mass unit of a substance by 1.00°C
Exothermic Reactions
If the products have stronger bonds than reactants, then the products have lower enthalpy than the reactants and are more stable; in this case, energy is released by the reaction, or the reaction is exothermic.
Spontaneous
Not requiring an outside source of energy to proceed
Anode
Where oxidation occurs and the solution is becoming more positively charged, the negative anions from the salt bridge solution flow into the half-cell.
Linear Geometry
- Central atom with 2 electron pairs - Zero lone pairs - sp hybridization - Ex. BeCl₂ and CO₂ B - A - B
Folded square, seesaw, distorted tetrahedron geometry
- Central atom has 5 electron pairs - One lone pair - dsp³ hybridization - Ex. SF₄, IF₄⁺
Solubility Rules
1. Compounds with an alkali metal cation (Na⁺, Li⁺, K⁺, etc) or an ammonium cation (NH₄⁺) are always soluble. 2. Compounds with a nitrate (NO₃⁻) anion are always soluble.
Bent Geometry
- Central atom has four electron pairs - two lone pair - sp³ hybridization - 109.5° bond angle - Ex. H₂O, OF₂, NH₂⁻
Amorphous Solid
Solids whose particles have no orderly pattern.
Dalton's Law
Ptotal = Pa + Pb + Pc + ...
Partial Pressure
Pa = (Ptotal)(Xa) Xa = moles of gas A/total moles of gas
Kinetic Molecular Theory
- The kinetic energy of an ideal gas is directly proportional to its absolute temperature: The greater the temperature, the greater the average kinetic energy of the gas molecules. - There are no forces of attraction between the gas molecules in an ideal gas. - Gas molecules are in constant motion, colliding with one another and with the walls of their container without losing any energy.
Oxidation
Electron loss
Solving Electroplating Problems
1. If you know the current and time, you can calculate the charge in coulombs. I = q/t I = current (amperes, A) q = charge (coulombs, C) t = time (seconds, s) 2. Once you know the charge in coulombs, you know how many electrons were involved in the reaction. moles of electrons = (coulombs) / (96,500 coulombs/mol) 3. When you know the number of moles of electrons and you know the half-reaction for the metal, you can find out how many moles of metal plated out. 4. Once you know the number of moles of the metal, you can convert this to the number of grams of the metal.
Linear Geometry
- Central atom has 5 electron pairs - Three lone pairs - dsp³ hybridization - Ex. XeF₂, I₃⁻
Boyle's Law (If temperature is constant)
As pressure increases, volume decreases, and vice versa
Trigonal Bipyramidal Geometry
- Central atom has 5 electron pairs - Zero lone pairs - dsp³ hybridization - Ex. PCl₅, PF₅
Trigonal Pyramidal Geometry
- Central atom has four electron pairs - One lone pair - sp³ hybridization - 109.5° bond angle - Ex. NH₃, PCl₃, AsH₃, SO₃²⁻
Enthalpy Change in Bonds
When bonds are broken, energy is released. When bonds are formed, energy is absorbed.
Charles' Law (If pressure is constant)
As temperature increases, volume increases
Square Pyramidal
- Central atom has 6 electron pairs - One lone pair - d²sp³ hybridization - Ex. BrF₅, IF₅
Octahedral Geometry
- Central atom has 6 electron pairs - Zero lone pairs - d²sp³ hybridization - Ex. SF₆
Section 3
(37 cards)