Chemistry Exam Review - ALL UNITS (2-6)

Describe atomic radius

The size of the atom

• Atomic radius INCREASES down a group 

  • More orbits = electrons are farther from nucleus (BIGGER ATOMIC RADIUS)

• Atomic radius DECREASES across a period

  • More protons = STRONGER PULL on ELECTRONS (SMALLER ATOMIC RADIUS)

Describe effective nuclear charge

• Effective Nuclear Charge (ENC) exerts a pull on the valence electrons

  • ENC = # protons - # core electrons

  • THINK “more protons = more ENC”

  • HIGHER ENC
    = STRONGER attraction of valence electrons
    = SMALLER atomic radius

Describe the ionic radius of cations

• Positive Ions (Cations)

  • LOSE electrons

  • Lose 1 energy level/shell

Radius DECREASES

Cations → LOSE electrons → Fewer orbits

= SMALLER ionic radius

Describe the ionic radius of anions

• Negative Ions (Anions)

  • GAIN electrons

  • Radius INCREASES due to electrons repelling each other (think: it needs to  “make room” for extra e^-)

  • Anions → GAIN electrons → Electron repulsion

    = LARGER ionic radius

Describe the Reactivity of METALS

• Metals react by LOSING electrons

  • Metal reactivity INCREASES down a group

    • Valence electrons FARTHER from the nucleus (more shielded) =  EASIER TO LOSE

  • Metal reactivity DECREASES across a period

    • # of valence electrons INCREASES = more valence electrons need to be given away (takes more energy to do this - HARDER TO LOSE) 

    • # of protons increases = MORE attraction of electrons (HARDER TO LOSE)

Describe the Reactivity of NONMETALS

• Nonmetals react by GAINING electrons

  • Nonmetal reactivity DECREASES down a group

    • Valence electrons are FARTHER from the nucleus (more shielded) = HARDER TO GAIN

  • Nonmetal reactivity INCREASES across a period

    • # of valence electrons INCREASES = fewer valence electrons need to be gained (takes less energy - EASIER TO GAIN) 

    • # of protons increases = MORE attraction of electrons (EASIER TO GAIN)

describe electron shielding

• outer (valence) electrons are partially shielded from the attractive force of the protons in the nucleus by the orbits in between the nucleus and the valence shell

  • More orbits = more shielding = e^- lost more easily

Describe ionization energy

• Ionization Energy (IE) - amount of energy required to remove an electron from the atom or ion (in gaseous state)

X_(g) + energy → X⁺ _(g) + e^-

• How hard it is for an atom to lose an electron (harder = ↑ IE)

Ionization energy INCREASES as you move left to right ACROSS a period

• MORE PROTONS
  = higher ENC
  (effective nuclear charge)
  = smaller atomic radius
  = MORE energy needed to remove an electron

• Ionization energy DECREASES as you move DOWN a group

• MORE ELECTRON SHELLS/ORBITS

= electrons farther from nucleus

= electron shielding

= LESS energy needed to remove an electron

Describe second ionization energy

• 2nd Ionization Energy (2nd IE) - amount of energy required to remove a SECOND electron from an ion (in gaseous state)

• It is always harder to remove a 2nd electron

• Removing 1st electron DECREASES
  atomic radius

  • Fewer electrons repelling each other & stronger pull from protons

  • Takes MORE ENERGY to remove a 2nd electron

Describe electron affinity

• lectron Affinity (EA) - The energy RELEASED from a (gaseous) atom ACCEPTING an electron 

How much an atom wants to gain an electron

• Inversely related to atomic radius

• Electron affinity INCREASES as you go left to right across a period

SMALLER atomic radius
= electrons closer to nucleus
= easier to gain/attract electrons

• Electron affinity DECREASES as you go down a group

  • MORE ELECTRON SHELLS/ORBITS
    = electrons farther from nucleus
    = more shielding
    = harder to gain/attract electrons

Describe electronegativity

• Electronegativity (EN) - the tendency for an element to attract shared electrons in a chemical bond

How strongly an element pulls the electrons to their side of the bond

• HIGHER EN means MORE electron attraction for an element

• Electronegativity INCREASES as you go left to right across a period

  • SMALLER atomic radius
    = electrons closer to nucleus
    = more pull on the electrons in a bonds
    
    Electronegativity DECREASES as you go down a group

    • MORE ELECTRON SHELLS/ORBITS
      = electrons farther from nucleus
      = more shielding
      = less pull on the electrons in a bond

Intramolecular vs. Intermolecular forces

• INTRAmolecular forces - attractive forces WITHIN a molecule

  • TIP - think “intramurals” (within a school)

  • Ionic, polar, or nonpolar
    

• INTERmolecular forces - attractive forces BETWEEN molecules

  • TIP - think “interact” (with others)

  • The stronger the intermolecular force,
    the higher the melting/boiling point
    

• The type of intramolecular force (ionic vs. covalent) determines
  the type of intermolecular force

Types of intermolecular forces

1. London Dispersion

   1. Weakest force

   2. Between ALL molecules (polar and nonpolar)

2. Dipole-Dipole

   1. Pretty strong

   2. Forces between ONLY polar molecules

   3. The greater ΔEN the stronger the dipole-dipole forces

3. Hydrogen Bonding

   1. Strongest intermolecular force

   2. Generally occurs when hydrogen bonds with N, O, or F (makes an extremely polar bond; very high ΔEN)

   3. REMEMBER - H=NOF

Mass of subatomic particles

Protons = 1

Neutrons = 1

Electrons = less than 1

What are the products of a decomposition reaction where a HYDRATE is the reactant?

Ionic compound + water (dehydration)

What are the products of a decomposition reaction where a METAL NITRATE is the reactant?

metal nitrite + oxygen

What are the products of a decomposition reaction where a METAL CARBONATE is the reactant?

Metal oxide + carbon dioxide

What are the products of a decomposition reaction where a METAL HYDROXIDE is the reactant?

Metal oxide + water

What does a nonmetal oxide + water form?

Acid

What does a metal oxide + water form?

Base

What are the prefixes and suffixes for derived acids?

Per ——- ic acid = 1 more oxygen

--ous = minus 1 oxygen

hypo —— ous = Minus 2 oxygen

Name the strong acids and bases

Acids - HI, HBr, HCl, H₂SO₄, HNO₃, HClO₃

Bases - Group 1 and 2

Boyle’s Law

Volume and temp inversely related

P₁V₁ = P₂V₂

Charle’s law

Volume and temperature are directly related

V₁ = V₂

T₁ T₂

Gay-Lussac’s Law

Pressure and temp are directly related

P₁ = P₂

T₁ T₂

Combined gas law

P₁V_{1 =} P₂V₂

T₁ T₂

Ideal Gas law

PV = nRT

Avogadro’s law

Equal volumes of gases have equal number or molecules and moles of molecules

V x mole ratio

Molar Volume

STP = 22.4 L/mol

SATP = 24.8 L/mol

n = V / V_m