Chem Multiple Choice Quiz 3

Standard temperature and pressure

• 0 °C (273 K) and 1 atm

• the molar volume at STP for an ideal gas is 1 mol gas = 22.4 L

Boyle’s Law

• P1​V1​=P2​V2​

• how the volume of a gas changes if the pressure is doubled 

Charles’ Law

• V1/T1 = V2/T2 

• predict what happens to the volume of a balloon if it's heated or
  cooled

Gay-Lussac’s Law

• P1/T1 = P2/T2

• what happens to the pressure of a sealed container if it is
  heated

Avogadro’s Law

• V1/n1 = V2/n2

• how volume changes as the amount (moles) of a gas changes

Combined gas law

When more than 1 variable is changed, how does that affect
the final values, which variable has a greater effect

Ideal gas law

• PV= nRT

• An ideal gas is a hypothetical substance that follows these rules perfectly

• Molecules have negligible volume compared to the container's volume

• There are no attractive or repulsive forces between the molecules

• Collisions are perfectly elastic, meaning no energy is lost

Relationship between Density and Temperature: a gas sample is heated in a
flexible container (at constant pressure) Volume increases and density
decreases

d ∝ 1/𝑇

Relationship between Density and Pressure in a sealed container that cannot
expand an increase of pressure on a gas increases the density

d ∝ P

Relationship between Density and Molar Mass: At the same temperature and
pressure, the gas that has a higher molar mass would have the higher density

d ∝ 𝑀

Explain the principle behind effusion/diffusion and how it relates to gas density
and molar mass

• Effusion and diffusion occur due to the random motion of gas particles

• Rate ∝ 1 / √(molar mass) → lighter gases spread or escape faster

• Higher molar mass or density = slower movement, lower molar mass or density = faster movement.

Partial pressures

the total pressure of a mixture of gases is the sum of their individual partial pressures

Grahams’ Law

the rate of diffusion or effusion of a gas is inversely proportional to the square root of its molar mass

What is the relationship between frequency and wavelength?

inverse

What is the relationship between frequency and energy?

direct

How can frequency be converted to wavelength

divide the speed of light by the frequency

How can frequency be converted to energy

multiply frequency by Planck’s constant (6.63 × 10 ^-34)

Bohr model of the Hydrogen atom and Energy of electron transitions

triangle E = - 2.28 x 10-18 (1/nf^2 – 1/ni^2)

Electrons orbit a nucleus in specific, allowed paths called

energy levels or orbits

Energy levels are quantized according to the equation

E = - 2.28 x 10-18 (1/n2)

Electrons can jump between these levels by absorbing or emitting

specific amounts of energy as photons

The emission spectra of elements like hydrogen result from

transitions from higher energy orbits to lower energy ones

A negative value indicates the photon being emitted while the positive value
indicates

the photon being absorbed

The further apart the orbits are

the larger the energy

Quantum numbers 

• a set of four numerical values that describe the unique location and properties of an electron in an atom

• they specify the electron’s energy level, shape of its orbital, orientation in space, and spin direction

Atomic orbitals

• regions around the nucleus of an atom where there is a high probability of finding an electron

• characterized by the first three quantum numbers (𝑛, 𝑙 and
  𝑚𝑙)

Principal Quantum Number (n)

• This number indicates the electron's main energy level,
  or shell, and can be any positive integer (1,2,3, . ..)

• As 𝑛 increases, the electron's average
  distance from the nucleus and its energy increases

Angular Momentum Quantum Number (𝑙)

• This number describes the shape of an atomic orbital and specifies the subshell

• 𝑙=0 corresponds to an s orbital (spherical)

• 𝑙=1 corresponds to a p orbital (dumbbell-shaped)

• 𝑙=2 corresponds to a d orbital (cloverleaf-shaped)

Magnetic Quantum Number (𝑚𝑙)

• This number describes the orientation of the orbital in
  space

• Its possible values range from −𝑙 to +𝑙, including zero

• For a given subshell, there are 2𝑙 + 1 possible orientations corresponding to the odd integers (1, 3, 5, 7)

Spin Quantum Number (𝑚𝑠)

• This number describes the intrinsic angular momentum, or
  "spin," of an electron

• There are only two possible values for 𝑚𝑠: + ½ or – ½ (“up” or
  “down” arrows)

Electron configuration 

• shows the arrangement of electrons in an atom’s orbitals

• tells you which orbitals are occupied and how many electrons are in each

How to write an electron configuration

1. Find the atomic number of the element.

• This equals the number of electrons in a neutral atom.

2. Use the Aufbau order (1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d → 7p) to assign electrons to orbitals.

3. Write the configuration by listing the orbitals in order with their electron counts as superscripts.

• Example: 1s² 2s² 2p⁶ 3s²

4. Check your total number of electrons — it should equal the atomic number

Aufbau Principle 

• Electrons fill the lowest energy orbitals first before moving to
  higher energy ones

• (1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d → 7p)

Pauli Exclusion Principle

• No two electrons in an atom can have the same set of
  four quantum numbers

• This means each orbital can hold a maximum of two
  electrons, and they must have opposite spins

Hund’s Rule

• Within a subshell, electrons will occupy each orbital individually
  before any orbital gets a second electron

• For example, in a p subshell, one electron will go into each of the three p orbitals before a second electron is added to any of them

Valence electrons 

• the outermost electrons of an atom

• they are responsible for an element's chemical properties and reactivity

• they are the electrons that are gained, lost, or shared during chemical reactions and used to create chemical bonds

How to determine how many valence electrons an atom or ion has

Atom

1. Find the element on the periodic table.

   • Look at its group number (column number).

Use the group number to find valence electrons:

Group	Elements	Valence Electrons
1 (IA)	H, Li, Na, K...	1
2 (IIA)	Be, Mg, Ca...	2
13 (IIIA)	B, Al, Ga...	3
14 (IVA)	C, Si, Ge...	4
15 (VA)	N, P, As...	5
16 (VIA)	O, S, Se...	6
17 (VIIA)	F, Cl, Br...	7
18 (VIIIA)	He, Ne, Ar...	8 (except He has 2)

3. Confirm by writing the electron configuration and identifying electrons in the outermost shell (highest n value).

Example:

• Oxygen: 1s² 2s² 2p⁴ → highest n = 2 → 2 + 4 = 6 valence electrons

• Sodium: 1s² 2s² 2p⁶ 3s¹ → highest n = 3 → 1 valence electron

Ion 

1. Start with the neutral atom’s valence electrons.

2. Add or remove electrons depending on the ion’s charge:

• For cations (positive ions): subtract electrons.

  • Example: Na → 1 valence e⁻ → Na⁺ → 0 valence electrons (empty outer shell; next lower shell is full).

• For anions (negative ions): add electrons.

  • Example: Cl → 7 valence e⁻ → Cl⁻ → 8 valence electrons (stable octet).

How to determine how many valence electrons there are in an atom
from electron configuration or position on the periodic table

Periodic Table

1. Find the element’s group number (the column it’s in).

2. Use the group number to determine valence electrons for main group elements (Groups 1–2 and 13–18).

Group	Example Elements	Valence Electrons
1 (IA)	H, Li, Na, K	1
2 (IIA)	Be, Mg, Ca	2
13 (IIIA)	B, Al, Ga	3
14 (IVA)	C, Si, Ge	4
15 (VA)	N, P, As	5
16 (VIA)	O, S, Se	6
17 (VIIA)	F, Cl, Br	7
18 (VIIIA)	He, Ne, Ar	8 (He has 2)

Electron configuration 

1. Write or look up the atom’s electron configuration.

2. Identify the highest principal energy level (n) — this is the outermost shell.

3. Count all the electrons in orbitals with that same n value — those are the valence electrons

Atoms form ions to achieve

a stable electron configuration

Metals typically have few valence electrons and lose them to become stable. This
results in a positive charge because

the atom now has more protons (positive) than
electrons (negative)

• the resulting positive ion is called a cation

Nonmetals have larger numbers of valence electrons (e.g., 6 or 7) and tend
to gain electrons. Gaining negatively charged electrons results in a negative ion which is called an

anion

Trends in metallic character

• Metals are defined as elements that readily lose their electrons to make cations

• Metallic character decreases from left to right across a period and increases
  from top to bottom down a group on the periodic table. This is because as you move across a period, the number of protons increases, pulling electrons closer and making them harder to lose

• As you move down a group, the atomic radius increases, placing outer electrons further from the nucleus, and shielding by inner electrons makes them easier to lose

Trends in radius size across a period or down a group

• Across a period (left to right) atomic radius decreases

• As you move across a period, the number of protons in the nucleus increases, creating a stronger positive charge

• Down a group (top to bottom) atomic radius increases 

Cations

• positive ions 

• smaller than their parent atoms because they have lost
  their valence shell

Anions

• negative ions

• larger than their parent atoms due to increased electron-electron repulsion

Ionization energy

• refers to the amount of energy needed to remove an electron from an atom

• across a period (left to right) ionization energy increases

• down a group (top to bottom) ionization energy decreases

Electron affinity

• the energy emitted when an atom takes in an electron

• across a period (left to right) electron affinity becomes more negative

• down a group (top to bottom) electron affinity becomes less negative

Electronegativity 

• the ability of an atom to attract shared electrons in a chemical bond 

• atoms with high electronegativity (like nonmetals) pull electrons more strongly 

• across a period (left to right) electronegativity increases 

• down a group (top to bottom) electronegativity decreases 

Where are the most electronegative elements found

the top right of the periodic table

Where are the least electronegative elements found

the bottom left of the periodic table

How to determine bond polarity

subtract the electronegativities of the 2 atoms in a bond to
determine if the bond is polar ( If the electronegativity difference < 0.4 is nonpolar, > 0.4
is polar – if it has a metal, it is ionic)

Covalent bonds

• the sharing of electron pairs between atoms

• it is created when a valence orbital from an element containing 1 electron overlaps with the valence orbital of another element also containing 1 electron

Lewis structures 

• diagrams that show how valence electrons are arranged in a molecule 

• bonds between atoms shown as lines and non-bonding electrons shown as dots 

• a single line is a single bond (2 electrons), a double line is a double bond (4
  electrons), and a triple line is a triple bond (6 electrons)

Formal charge

• the hypothetical charge an atom would have if bonding electrons were split equally between atoms and the molecule's total charge was distributed equally

• the number of valence electrons for a particular atom minus the number of
  assigned electrons determines the formal charge on the atom

Resonance

• a way to describe the positions of electrons in certain molecules or
  ions where a single Lewis structure is not enough to represent the true
  bonding

• this lowers the energy of the molecule, making it more stable than one particular structure

• to identify molecules with resonance look for a multiple bond next to an atom
  that is capable of taking part of that multiple bond, usually on that has a lone
  pair of electrons

VSEPR theory 

• a model used to predict the 3D shape of a molecule

• To use VSEPR, draw a Lewis structure, count the total number of electron group: bonds (single bonds, double bonds, triple bonds are all 1 bond) and lone pairs around the central atom, and determine the geometry that positions these groups at maximum distance from each other