U2

solutions

• solute dissolved in solvent

• indistinguishable (homogenous)

• dissolved particles are too small to see

homogenous

• uniform

• well mixed

• miscible

• eg; NaCl in water

heterogenous

• not uniform

• not well mixed

• immiscible

• eg; choc chip cookie

aqueous

solid / liquid / gas dissolved in water

unsaturated

more solute can be dissolved in solution

saturated

no more solute can be dissolved at current temp + pressure

supersaturated

• a solution containing more than the max solute that can be dissolved at given temp + pressure

• this is done by increasing temp —> adding solute —> decreasing temp

suspensions

• heterogenous mixture

• solute not dissolved significantly in a solute

• some particles settle over time + can be separated

• eg; RBC’s in plasma

colloids

• WEIRD LIQUIDS

• mixture of particles consisting of smaller clusters of ions + molecules

• evenly dispersed throughout the solvent + don’t settle over time

• eg; mayo, ink, milk

dissolution

• solute + solvent particles attract each other strongly enough to overcome the forces holding their own particles together

• this allows the solute to disperse throughout the solvent

solubility formula

m (solute in grams)

—————————

100g water

• most solids increase solubility as temp increases

• most gases decrease solubility as temp increases

• solubility of a substance can be graphed on a solubility curve

electrolyte

solution formed when solute dissolves to form ions

strong electrolyte

• all of the solute dissolves to form ions

• ionic compounds fully dissociate to form strong electrolytes

• strong acids fully ionise to form strong electrolytes

weak electrolyte

• some of the solute dissolves to form molecules - no ions formed

• weak acids partially ionise to form weak electrolytes

like dissolves like

• polar solvents dissolves polar molecules

• non-polar solvents dissolves non-polar molecules

precipitation reactions

• occurs between cations + anions in aqueous solutions

• combine to form an insoluble ionic solid (precipitate)

• eg; potassium carbonate + copper (II) nitrate

  full: K₂CO_3(aq) + Cu(NO₃)_2(aq) —> 2KNO_3(aq) + CuCO_3(s)

  ionic: CO₃^-2_(aq) + Cu²⁺_(aq) —> CuCO_3(s)

  observations: a colourless and blue solution is mixed together + a green precipitate if formed

concentration

amount of solute + solvent present in solution

n = c v

moles of solute = conc.(mol / L) x volume (L)

C1V1 = C2V2

initial = dilute

properties of acids

• turns litmus paper red

• corrosive

• sour

• solutions have pH <7

• solutions conduct electrical current

properties of bases

• turns litmus paper blue

• slippery

• bitter

• solutions have pH >7

• solutions conduct electrical current

strong acids

• completely ionises in water

• HCl, HNO₃, H₂SO₄

weak acids

• only partially ionises in water

• CH₃COOH, H₂CO₃, H₃PO₄

bases

• only partially ionises in water

• NH₃, metal hydroxides, metal oxides

acid reactions

Acid + Metal —> Salt + Hydrogen Gas

Acid + Metal Oxide —> Salt + Water

Acid + Metal Hydroxide —> Salt + Water

Acid + Metal Carbonate—> Salt + Water + Carbon Dioxide

Acid + Metal Hydrogen Carbonate—> Salt + Water + Carbon Dioxide

Acid + Metal Sulphite —> Salt + Water + Sulphur Dioxide

Base + Ammonium Salt —> Salt + Water + Ammonia

Base + Non-Metal Oxide —> Salt + Water

Arrhenius theory

• acids produce H⁺ ions when they dissolve in water

• bases produce OH^- ions when they dissolve in water

• doesn’t account for all acid-base behaviour

  • eg; why can HCO₃^- ions act as an acid + a base?

• can only be used when talking about aqueous solutions

Bronsted-lowry theory

• acids will donate protons (H⁺ ions)

• bases will accept protons (H⁺ ions)

beaker

conjugate acids + bases

• once a base has accepted a proton, it has the potential to donate the proton (act as a base)

• once an acid has donated a proton it has the ability to gain it back (act like a base)

indicators

• acid and base react with indicators

• changes their colours

• allows the pH of the solution to be determined

• litmus paper: red = acidic blue = basic

• universal indicator: changes colour over the whole pH scale

pH

• describes the conc of H⁺ ions in solutions

• pH = -log [H⁺]

• [H⁺] = 10^-pH

• 0 = acidic

• 7 = neutral

• 14 = basic

monoprotic acids

• can make 1H⁺ per molecule

polyprotic acids

• can make 2 or more H⁺ per molecule

electronegativity (EN)

• ability to attract a bonding pair of electrons to form a covalent bond

• increases left—> right

  • higher nuclear charge

• increases down —> up

  • less electron shells

polar

• uneven distribution of charge

• all ionic compounds are polar

• for covalent molecules, the polarity is determined by the shape + direction of polar bonds (if present)

polar bonds

• 2 different elements are sharing electrons with different EN (ability to attract / bond with electrons)

• 2 of the same elements are non-polar since they have the same EN

big dipole movement

very different EN

small dipole movement

very similar EN

VSEPR theory

• the valence shell electron pair repulsion theory

• valence electrons repel each other

• they will spread apart as much as possible to minimise the repulsion between them.

linear

• 1 or 2 bonds

• no lone pairs

• 180 degrees

• non - polar: both of the atoms coming off the central atom are the same element (equal charges in opposite directions cancel each other out

  • eg; CO₂

• polar: not all the same element

bent

• 2 bonds

• 1 or 2 lone pairs

• v shaped

• polar: not all the same element

trigonal planar

• 3 bonds

• no lone pairs

• 120 degrees

• non-polar: all atoms coming off the central atom are the same

  • eg; BF₃

• polar: not all the same element

pyramidal

• 3 bonds

• 1 lone pair

• polar: not all the same element

tetrahedral

• 4 bonds

• no lone pairs

• 105.5 degrees

• non-polar: all atoms coming off central atom are the same

  • eg; CH₄

• polar: not all the same element

shape of molecules can determine physical properties of covalent molecular substances, such as:

• vapour pressure

• melting point

• boiling point

• solubility

shape of molecules determines…

…how it will interact with other molecules

lone pairs

electrons around central atom that are not involved in bonding

head to tail

draw dipoles with head touching tail

if its a closed shape it is non-polar since charges cancel out

if it is an open shape is is polar since charges do not cancel out

IMF’s

• electrostatic force

• between the (+) and (-) charges in the molecules as a result of uneven distribution of electrons

• increase EN = more likely an atom in a molecule will attract electrons towards it and away from atom with lower EN

polar diatomic molecules

• 2 different elements with different EN in a covalent bond are polar

  • electrons will stay closer to the most EN atom as it has the stronger pull on electrons.

non-polar diatomic molecules

• 2 of the same atoms with same EN in a covalent bond are non-polar

  • electrons are shared equally

  • no charge on either end of the molecule.

polar polyatomic molecules

• Molecules made up of more than two atoms

• Asymmetrical molecules are polar

non-polar polyatomic molecules

• Molecules made up of more than two atoms

• Symmetrical molecules are non-polar

3 types of IMF forces

1. dispersion forces

2. dipole-dipole

3. hydrogen bonding

dispersion forces

• all molecules

• electrons are constantly moving

• at any given time there will be more on one side of the molecule than the other

• causes a temporary dipole

• increase dispersion forces —> increase size of molecule

  • more electrons = more temporary dipoles

• dispersion forces is the weakest IMF

dipole-dipole

• polar molecules

• dipole on one molecule is attracted to the oppositely charged dipole on neighbouring molecule

• stronger than dispersion forces as dipole is not temporary

hydrogen bonding

• a H bonded to a N,O or F (most EN elements)

• lone pairs on the N,O or F

• attraction between the (+)H and lone pair of electrons

• strongest IMF