structure and bonding

ionic bond

bond between positively and negatively charged ion with strong attraction. uses electrostatic attraction. ionic compounds must have a net charge of zero

cation/anion

cation - positively charged (lose electrons)

anion - negatively charged (gain electrons)

common polyatomic cations/anions to remember

Ammonium NH4+

Hydroxide OH-

Carbonate CO32-

Sulfate SO42-

Nitrate NO3-

covalent bonds

formed between non metals sharing valence electrons to achieve a full outer shell

inter vs intra molecular bonds

inter → between two molecules

intra → within a molecule

lewis diagram

• the atom forming the most bonds is the center

• all atoms should have 8 valence electrons

• lewis diagrams determine molecular 3D shapes

electron density

pairs of electrons around the central atom. electron density can be bonding and non bonding.

note: double and triple bonds are only counted as one area of electron density

VSEPR (valence shell electron pair repulsion) theory

electron pairs are negatively charged, so one pair repels another. electron pairs will always be the furthest distance apart from each other to reduce repulsion

carbon typical arrangement

normally four single bonds but in the case of CO₂, it is different

electronegativity

atoms ability to attract the bonding electron to itself. electronegativity in covalent bonds = sharing of electrons may or may not be equal depending on the electronegativity of the atoms involved

electronegativity trend

from groups 1 to 17, electronegativity values increase and as you go up the columns, the electronegativity increases. fluorine is the most electronegative of the elements (with a maximum number of 4 - technically 3.98) ELECTRONEGATIVITY VALUES HAVE NO UNITS

why = increasing nuclear charge due to increase in number of protons which increases nuclear charge and attraction to valence shell electrons.

→ electronegativity decreases down groups as atomic radius get larger. valence electrons get further away, reducing atoms ability to attract electrons

non polar covalent bonds

atoms share electrons equally WHEN atoms have identical electronegativity values OR very little difference (<0.4 difference)

polar covalent bonds

where there is a significant difference in the electronegativity of the two atoms (between 0.4 and 1.7) shared electrons are located towards the atom that has the higher electronegativity number because it has the stronger ability to attract the bonding electrons. there are two ends, slightly negative (more electrons) and slightly positive (less electrons)

dipole movement

has both direction and magnitude. always points to the electron rich side. may have multiple polar covalent bonds present. dipole movements can cancel out and reinforce each other. DEAW ARROW TO SHOW DIRECTION

linear

when there are only two atoms bonded together (e.g. HCl)

OR when the central atom is surrounded by only two areas of electron density (both areas are bonding pairs) (e.g. CO₂ or BeH₂)

bond angle = 180^o

trigonal planar

when the central atom is surrounded by 3 areas of electron density (all bonding pairs) e.g. BF₃

bond angle = 120^o

tetrahedral

when the central atom is surrounded by 4 areas of electron density (all bonding pairs) e.g. CH₄

bond angle = 109.5^o

bent

a central atom bonded to two other atoms and has one non bonding pair of electrons (3 areas of electron density but only two bonding pairs) e.g. O₃

bond angle = <120^o or 117.5^o

trigonal pyramid

a central atom bonded to three other atoms but has a non bonding pair of electrons (4 areas of electron density but only three bonding pairs) e.g. NH₃

bond angle = 109.5^o or 107^o

bent (but with four regions)

a central atom bonded to two other atoms with two non-bonding pairs of electrons (4 areas of electron density but only two bonding pairs) e.g. H₂O

bond angle = 109.5^o or 104^o

covalent networks

large networks of atoms held together with covalent bonds. e.g. carbon dioxide or silicon dioxide

2D covalent networks - graphite (allotrope of carbon)

carbon has 4 valence electrons but only forms three strong bonds - leaving a electron free. → arranged in flat sheets as hexagonal rings

3D covalent networks - diamond, allotrope of carbon

carbon is bonded to 4 other carbon atoms. forms a tetrahedral arrangement

3D covalent networks - silicon dioxide

silicon is bonded to four other oxygen atoms - each oxygen atom is bonded to two silicons

hardness + brittleness of ionic solids

ionic solids are very hard as they are closely packed due to ionic bonds

ionic solids are brittle (shatter with force) as it causes positive ions to line up with other positive ions (same with negative ions) and the like charges repel strongly

forces in molecular substances

molecular substances have forces between particles NOT covalent bonds

in molecular solids, the attractive forces between particles are weak, intermolecular forces

molecular mass

all the molecules together

melting and boiling point of INTERMOLECULAR (weak) attractive forces

generally low melting and boiling points. number of intermolecular forces and relative strength increases with size of molecular mass. the more intermolecular forces there are & the bigger the molecule, the higher the melting and boiling point

WHEN they are in a SOLID state, they are brittle → it is very easy to break the intermolecular bonds (molecular substances tend to be liquid or gas at room temperature)

covalent network hardness

covalent networks held together by strong covalent bonds = high melting point

2D - graphite is soft as weakly bonded layers

3D - diamond is one of the hardest substances because of tetrahedral arrangement

metal melting points/hardness

metals have high melting points generally

metallic bond is very strong

typically at room temp, metals are solid

malleable and ductile → atoms slim past each other and remain bonded

NEEDED FOR EXCELLENCE

COMPARE AND CONTRAST TWO DIFFERENT SUBSTANCES IN TERMS OF STRUCTURE, BONDING AND PROPERTIES

solubility

ability of substance (solute) to dissolve into a specific solvent

for the substance to be soluble, it needs to have a greater attraction to the solvent rather than itself in order to dissolve

ionic solids solubility rules

• soluble in polar substances

• ions are charged particles and strong electrostatic attraction

• polar solvent is able to overcome electrostatic attraction, breaking apart lattice - causing it to dissolve

HOWEVER

• they are insoluble in non-polar solvents

• non-polar solvents do not have a strong attraction with ions, cannot overcome electrostatic attraction

metal solids solubility rules

insoluble in any kind of substance, metallic bond is too strong to overcome

covalent networks - solubility rules

insoluble in any type of solvent - covalent bond is too strong to overcome

molecular substance solubility

depends on relative strength of the intermolecular forces between molecules and the attractive force in the solvent molecule

• if the molecule structure is polar, it will be more attracted to a polar solvent that its own molecules, so it dissolves in polar substances

• if the molecule structure is non polar, it will be less attracted to a polar solvent that its own molecules so it will not dissolve in polar substances

• if the molecule structure is polar, it will be more attracted to itself that non polar molecules and will not dissolve in a non polar substance

• if the molecular substance is non polar, it will be more attracted to a non polar solvent than its own molecules so it will dissolve in a non polar substance

endothermic

absorbs heat and has a positive Δ_rH. Energy is absorbed from surroundings into system. products have more energy than reactants. air around becomes cold/decrease in temperature

bond breaking = energy in = endothermic

exothermic

energy releases heat and has a negative Δ_rH. energy is lost to surroundings and exits system. products have less energy that reactants. temperature increases in the surroundings

bond making = energy out = exothermic

enthalpy

chemical potential stored within bonds. measured in Joules or kiloJoules

measure enthalpy when there is a chemical or physical change

Δ = change H = enthalpy

heat enthalpy equation

fancy symbol means sum of

bond enthalpy

the amount of energy that is needed to break the covalent bonds in one mole of a molecule in its gaseous form

bond enthalpy will ALWAYS be positive

double and triple bonds need more energy to be broken

CHEMICAL REACTIONS INVOLVE THE SIMULTANEOUS BREAKING OF EXISTING BONDS AND MAKING OF NEW BONDS