General Chemistry

ion

an atom that has gained or lost one or more electrons, resulting in a net electric charge

cation

a positively charged ion that results from the loss of one or more electrons from an atom

anion

a negatively charged ion that results from the gain of one or more electrons from an atom

compounds

molecules comprised of two or more different elements bonded together

(ex: H₂O, CO₂)

ionic compound

metal + non-metal (contain ionic bonds)

(ex: NaCl, CuCl₂)

molecular compounds

two or more non-metals (contain covalent bonds)

(ex: H₂O, CO₂)

metal, non metal, metalloid

blue = ___

red = ___

green = ___

polyatomic ions

make ionic compounds even though they do NOT contain any metals

(ex: ammonium chloride NH₄Cl. this is because NH₄⁺ is a cation and Cl^- is an anion)

NH₄⁺

ammonium

(polyatomic ion)

covalent bond

between non-metals

(atoms share their electrons, so they do NOT have true + or - charges. they can have partial + and - changes if they share their electrons unevenly)

(ex: oxygen hogs more e^- so hydrogens are partially + and oxygen is partially -)

increases

electronegativity ___ as you go up and to the right on the periodic table

polar covalent bonds

bonds in molecular compounds that have an uneven sharing of electrons and partial positive or negative charges

(caused by a significant electronegativity difference between the bonded atoms)

nonpolar covalent bonds

occur when nonmetal bonded atoms do NOT have a significant electronegativity difference between them

ionic compounds

• high melting points

• high boiling points

• brittle

• hard

(ex: NaCl, MgO)

molecular compounds

• low melting points

• do not conduct electricity

• intermolecular forces

(ex: H₂O, Cl₂)

metallic compounds

• variable hardness and melting points

• conduct electricity and heat

• lustrous (shiny)

• malleable

• ductile

metallic bonding

(ex: Fe, Mg)

network covalent compounds

• high melting points

• high boiling points

• hard

• do not conduct electricity

network of covalent bonds

(ex: C - diamond, graphite, SiO₂ - quartz)

lattice energy

energy required to completely separate an ionic compounds cations from its anions

bigger charges = larger ___ energy

shorter bond distance = larger ___ energy

smaller

atom sizes get ___ as you go to the right across a row or up a column on the periodic table

how many valence electrons?

• usually the same as its column number on the periodic table

• if d-block is completely filled, d-block electrons do NOT count

• if d-block is NOT completely filled, d-block electrons DO count

(ex: Zn is in the last row of the d-block, also known as fully filled.) Only count the 2 s-block columns, skip all of the d-block electrons → Zn has 2 valence electrons)

valence electron exceptions to full octet rule

• H: only needs 2 electrons, not 8

• Be: only needs 4 electrons, not 8

• B and Al: sometimes only have 6 electrons

electron configuration exceptions

• Cr

• Mo

• Cu

• Ag

• Au

take away one electron from s orbital and add to d orbital

paramagnetic

has unpaired electrons

paramagnetic elements are attracted to magnets

if an element has an ODD number of electrons → paramagnetic

“unpaired”-a-magnetic

diamagnetic

all electrons are paired

diamagnetic elements are slightly repelled by magnetic fields

if an element has an EVEN number of electrons → paramagnetic OR diamagnetic (fill out electron configuration energy diagram to find out)

“DI = 2 → electrons are paired (2) → repel”

electromagnetic spectrum (lowest energy / frequency to highest)

Radio waves → Micro-waves → Infrared radiation → Visible light → Ultraviolet → X-rays → Gamma-rays

“Roman Men Invented Very Unusual X-ray G*ns”

linear

• 2 electron groups (0 lone pairs)

• bond angle: 180°

trigonal planar

• 3 electron groups (0 lone pairs)

• bond angle: 120°

bent

• 3 electron groups (1 lone pair)

  • bond angle: <120°

OR

• 4 electron groups (2 lone pairs)

  • bond angle: <109.5°

diatomic molecules

“Have No Fear Of Ice Cold Beer”

H₂, N₂, F₂, O₂, I₂, Cl₂, Br₂

7 molecules total, form a 7 on the periodic table

• Hydrogen

• Nitrogen

• Fluorine

• Oxygen

• Iodine

• Chlorine

• Bromine

atomic radius of an ion is affected by

1. Number of electron shells: More electron shells → larger atomic radius because the electrons are spread out over a larger physical space.

2. Nuclear charge: A higher effective nuclear charge (more protons in the nucleus) → smaller atomic radius because the electrons are pulled closer to the nucleus.

electron affinity

the amount of energy released when an atom gains an electron

increases moving up and to the right on the periodic table

ionization energy

the amount of energy required to remove an electron

electronegativity

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

lattice energy

the energy required to separate an ionic compound’s cations and anions

effective nuclear charge

the force that attracts electrons toward the nucleus

first ionization energy

the energy required to remove the outermost electron from a neutral atom.

ionization energy increases moving up and to the right on the periodic table.

atomic radius

increases moving down and to the left on the periodic table

• this is due to the addition of electron shells and a decrease in effective nuclear charge