Chem - Ch 6

chemical bond

link between atoms that results from the mutual attraction of one atom’s nucleus for another atom’s electrons (and vice versa) resulting in going into a lower energy and more stable

ionic bond

transfer of electrons

covalent bonds

sharing of electrons

metallic bonds

happens in metals when there is only ONE TYPE of electrons (ex. Fe + Fe)

(covalent) polar bond

unequal sharing of electrons, more electronegativity on one side over the other (ex. HCl)

(covalent) non polar bonds

equal sharing of electrons

electronegativity difference - ionic

greater than 1.7 (metal+metal/polyatomic ion)

electronegativity difference - polar covalent

1.7-0.4 (two different nonmetals)

electronegativity difference - nonpolar

0.3 or less (two of the same nonmetals)

endothermic reaction

breaking bonds, must put energy in to break bonds

exothermic reaction

forming bonds, release energy when those bonds are broken

ionic compound

a compound made from positive and negative ions so that the charges are equal. it involves a transfer of electrons

formula unit

the lowest whole number ratio of ion in an ionic compound

metallic ionic bond

lose electrons due to low ionization energy

nonmetallic ionic bond

gain electrons due to high electronegativity

metals and nonmetals go to ___ to have 8 valence electrons

s²p^6

octet rule

the tendency to arrange electrons so each atom has 8

lattice energy

energy released when an ionic compound forms

bond strength relies on

the nuclear charge (number of protons) and the number of valence electrons

sea of electrons theory

the valence Electrons form a free moving sea of electrons that are attracted to multiple positive nuclei

no atoms want the electrons in metallic bonds, so they are

mobile

molecule

smallest quantity of matter that exists by itself and retains the properties of that substance

monatomic molecules

noble gases, He, Ne, Ar, Fe

diatomic molecules

H2, O2, N2, Cl2, Br2, I2, F2 (honcl brif!)

polyatomic molecules

P4, S8, etc.

bond length vs bond energy

bond energy goes up as bond length goes down

• shorter bonds are more stable and require more energy to break

• longer bonds are less stable and easier to break

• multiple bonds result in shorter bond lengths and higher bond energies

sigma bond

a single bond with head on head overlap of orbitals

pi bond

the second bond of a double bond with side to side double overlap of p orbitals (can also be the third bond of a triple bond)

Lewis Dot Diagrams

covalent compounds and polyatomic ions

1. each atom gets 8 electrons (except H gets 2)

2. each atom goes for close to the right number of bond

3. the atom that makes the most bonds goes in the middle (H is always on the outside)

4. the atom that you have the least of goes in the middle

5. the least electronegative atom is placed in the middle

   then check bonds —> check electrons

charges in Lewis dot diagrams

when atoms get less bonds than atoms normally get = negative charge

when atoms get more bonds than atoms normally get = positive charge

coordinate covalent bond

a bond where two shared electrons in a bond are donated by 1 atom

adding a hydrogen to the single bonded oxygen changes the charge to neutral

H’s go on O’s yeah

Bad Boys of the octet rule

Boron and Beryllium

RHED

bonds + lone pairs

any type of bonds

1 shared pair

linear (1 RHED)

• ex. AB

• no lone pairs

• 180 degree angle

linear (2 RHED)

• ex. AB2

• 2 shared pairs, no lone pairs

• 180 degree angle

bent (3 RHED)

• ex. AB2E

• 2 shared pairs, 1 lone pair

• 120 degree angle

bent (4 RHED)

• ex. AB2E2 (H2O)

• 2 shared pairs, 1 lone pair

• 104.5 degree angle

triangular planar (trigonal) (3 RHED)

• ex. AB3

• 3 shared pairs

• 120 degree angle

triangular pyramidal (pyramid) (4 RHED)

• ex. AB3E

• 3 shared pairs, 1 lone pair

• 107 degree angle

tetrahedron (4 RHED)

• ex. AB4

• 4 shared pairs

• 109.5 degree angle

triangular bipyramidal (5 RHED)

• ex. AB5

• 5 shared pairs

• 120 and 90 degree angles

octahedron (6 RHED)

• ex. AB6

• 6 shared pairs

• 90 degree angle

intermolecular forces (IMF)

forces that hold diatomic, monatomic molecules, and covalent bonds together

• determines the phase of the substance and some properties

intramolecular forces

chemical bonds (ionic, covalent, metallic)

• happens within a molecule

• always strong

dipole-dipole

forms when electrons are unevenly distributed - polar

hydrogen bonding

super duper dipole-dipole, only happens with FON

• this is what causes freezing water to expand

• super high boiling point, heat of vaporization, and melting point

• low vapor pressure

Van der Waals forces (London forces)

temporarily induced dipoles caused by the motion of electrons

• more electrons = more attraction

• nonpolar

bond energy

the energy needed to break a bond - measured in kJ/mole

• stronger bond = more stable = needs more energy to break the bond

• weaker bond = less stable = little energy to break the bond