Chem Unit 2 - IMF, Solutions and Reactions

NOT FINISHED, use equations when answering questions

VSEPR Theory

states that repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible

Steric Number

number of atoms bonded to central atom + number of lone pairs on central atom

Core Charge

A measure of the attractive force felt by the valence shell electrons towards the nucleus

Polar Bonds

a type of covalent bond between atoms that differ in electronegativity. the shared electrons are pulled closer to the more electronegative atom. making one slightly negative and the other slightly positive.

Dipole

permenant build up of negative charge at one end and positive charge at the other end of a covalent bond or molecule

Strength of IMF depends on

- shape and size of the molecule
- elements it contains

Dispersion Forces

attractions between molecules caused by the electron motion on one molecule affecting the electron motion on the other through electrical forces; these are the weakest interactions between molecules.

Predicting Strength of Dispersion Forces

- for larger atoms with a greater number of electrons, there is a greater chance that these electrons will be arranged asymmetrically at any instant, compared to the smaller number of electrons in smaller atoms.
- more electrons and bigger atomic radius means a higher boiling point
- shape also influences strength. larger the surface area over which electrons can develop an instantaneous dipole, and the closer molecules can fit next to one another, the stronger the dispersion force

Hydrogen Bonding

the directional intermolecular force in which a hydrogen atom that is bonded to either Nitrogen, Oxygen or Flourine is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule. covalent bond between these three atoms and hydrogen will be very polar and atomic radius of these atoms is small so dipoles are confined to a small space

Dipole Dipole Forces

electrostatic attraction between the partial negative and partial positive charges induced by permenant dipoles of atoms of a molecule.

Chromatography

- name given to the group of techniques that separate substances based upon differential distribution between a stationary phase and mobile phase. Relies on the chemical property that different substances will adsorb onto a surface and desorb into a solvent at different rates.

Thin Layer Chromatography (TLC)

- used to analyse the presence of particular drugs and amino acids
- faster and provides more separation then paper chromatography
- separation achieved by solvent (mobile phase) moving up the stationary phase (due to capillary action)
- goal is to measure the distance that each component has travelled from the starting line and compare this to the distance the solvent has advanced, called the Retention Factor (Rf)
- if a substance does not advance at all, Rf \= 0, if it moves at the same rate as the solvent, Rf \= 1
- Rf values can be compared to standards (pure samples run under same conditions)

Gas Chromatography (GC)

- separation technique for small organic molecules. based on volatility
- consists of a gas bottle, oven, column, detector and recorder
- sample is injected into oven where gas pushes sample into long, thin column
- larger particle that adsorb more readily take longer to leave column
- time sample takes to elute is characteristic for substance, called the retention time, or Rt
- technique can also identify amount of analyte
- data can be represented on a calibration curve
- peak area is in proportion to height of peak

High Performance Liquid Chromatography (HPLC)

- used for larger organic molecules
- can be used for substances that are unstable to heat because there is no oven
- stationary column shorter then gas chromatograph. Mobile phase is liquid not gas.
- Pump is used because pressure is requred to move liquid through the column
- Rt values are also used, as well as standards
- can be also used for quantitative analysis

Factors affecting retention time in GC

- volatility: more volatile (lower BP), more likely it will be in vapour phase and lower retention time
- column temp: higher this is, lower retention time
- carrier gas flow rate: greater this is, lower Rt

Solution

- A homogenous mixture where solute is dissolved in a solvent
- solutions can be solid, liquid or gas
- solute dissolved in water is called aqueous solution
- most solids are increasing solubility with increasing temperature
- gases are always less soluble at higher temperatures
- aqueous solutions are sometimes referred to as electrolytes

Ionic Compounds as Electrolytes

- all ionic compounds are strong electrolytes
- ionic compounds completely dissociate in solution, forming free mobile ions

Covalent Molecular Compounds As Electrolytes

- non-electrolytes (do not produce ions when dissolved in water)
- C6H12O6 (s) -\> C6H12O6 (aq)
- When these substances dissolve in water they produce independent free molecules only
- exceptions: covalent molecules that have acidic or basic properties

Acids and Bases as Electrolytes

- strong acids completely ionise to form strong electrolytes in solution of free mobile ions
- strong acids are HCl, H2SO4, HNO3
- weak acids or bases form weak electrolytes, dissolve in solution but go on to form some ions by process of ionisation (has equilibrium arrow)

Electrical Conductivity of Electrolytes

- electrical conductivity of a solution depends on ability of positive ions to move freely towards negative electrode, and negative ions to move freely towards positive electrode
- higher concentration of ions, the greater its ability to conduct electricity

Formation of Solutions

- in a homogenous mixture, there is only one phase
- if an insoluble substance like sand or oil is added to water, a heterogenous mixture is formed
- formation of solution involves the rearrangement of bonds.
- bonds between the particles of the solid (solute) are broken as particles move away to mix with the liquid
- bonds between the liquid are disrupted as solute particles move in between them and new bonds are formed between solute and solvent.
- energy is required to break a bond, the stronger the bond the more energy is required
- energy is released when a bond is formed, the stronger the bond formed the more energy is released
- if bonds being broken are of similar strength or weaker then new bonds being formed are of similar strength or weaker then new bonds being formed, then solute should dissolve the solvent.

Predicting Solubility

- Polar and Ionic Substances will dissolve in polar solvents
- Non-polar solutes will dissolve in non polar solvents
- solubility of polar solute will decrease as polarity of solvent decreases

Dissolving Ionic Compounds in Water

- attractive forces exist between ions and polar molecules such as water. called an ion dipole force and it can be stronger then hydrogen bonds

Spectator Ions

Ions that do not take part in a chemical reaction and are found in solution both before and after the reaction

Observations

following should be included:

• appearance of original reactants

• colour of solutions

• colour of precipitates

• whether a solid dissolves or is produced

• when there is a gas produced (includes colour and smell)

3 types of bonding in water

- strong covalent bonding between H and O within molecules
- hydrogen bonding (unique properties of water)
- dispersion forces which occur between molecules of water

Unique Properties of Water

Universal Solvent, Adhesion, Cohesion, high specific heat, density (expansion upon freezing), high BP and MP, surface tension, amphoteric

Surface Tension

- tendency of a liquid to resist any increase in surface area.
- water molecules attract one another as each molecule forms a bond with ones in its vicinity
- at surface, outermost layer of molecules has fewer molecules to cling to, and therefore compensates by estabilishing stronger bonds with neighbours, leading to formation of surface tension

Surfactant

any substance that interferes with the hydrogen bonding between water molecules and thereby reduces surface tension. e.g soap and detergents break hydrogen bonds and water spreads out rather then remain in a droplet

Cohesion

an attraction between molecules of the same substance.
- cohesive nature of water helps plants take up nutrients from soil
- capillary action relies on adhesion of water molecules to size of zylem vessel and cohesive properties of water
- cohesion of water molecules mean dissolved nutrients are pulled to the top of trees, defying gravity

Capillary Action

- due to 3 main forces:
- cohesive force: intermolecular forces between molecules in a substance that helps to maintain certain shape of a liquid
- surface tension: due to cohesive forces at surface of material, results in surface of fluid being under tension
- adhesive force: these are forces of attraction between unlike molecules

Multistage Flash Distillation

- sea water enters brine heater, heated seawater flows to first stages where it 'flashes' (vapourises) upon entry
- flashed vapour is condensed on outside of tubes carrying seawater and condensed steam withdrawn as fresh water
- unflashed portion of sea water is sent to second stage for further flashing (second stage) operates lower pressure to lower BP of sea water and more vapour is flashed off
- water is extracted and sea water is fed to third stage

Osmosis

- natural process by which water flows into plant and animal cells, through cellular membranes, AKA semi permeable memrbanes. solvents naturally move from area of low to high conc

Reverse Osmosis

- if external pressure is applied to cell with higher conc, pure water will be forced out of cell and salt remains behind SPM
- impurities are filtered incl. soil, microorganisms, suspended solids (colloids)
- highly efficient, easy to operate

Groundwater Remediation

1. Aeration: dissolved gases like CO2 and H2S removed by spraying water into air and oxygenated and chlorinated
2. Clariification: coagulate (alum and lime) is added to encourage floculation of suspended particles
3. Sand Filtration
4. Disinfection: chlorination and chloramination is used to treat pathogens in long network of pipes
5. Fluoridation: fluorine levels of 0.6-1.0 mg per Litre are maintained in drinking water supplies
6. pH: maintained between 6.5-8.5 pH by addition of lime, CaO or by dissolving CO2.

Properties of Acids

result from production of H+ ions (or H3O+ in solution):
- turn blue litmus red
- neutralised by alkalis (base)
- taste sour
- corrosive
- molecular in structure and dissolve in water to produce electrolyte

Salt

An ionic compound made from the neutralisation of an acid with a base.

Acid + reactive metal

salt + hydrogen gas (not acid/base reaction)

Acid + metal hydroxide

salt + water

acid + metal oxide

salt + water

acid + metal carbonate

salt + water + carbon dioxide

Acid + metal hydrogen carbonate

salt + water + carbon dioxide

acid + metal sulphite (e.g Na2SO3)

salt + sulphur dioxide + water

Properties of Bases

result from production of OH- ions in solution:
- turn red litmus blue
- are generally ionic (metal oxides/hydroxides) and dissolve in water to produce electrolytes
- neutralised by acids
- taste bitter
- feel slippery
- may be corrosive (caustic)

Acid + base

salt + water

Ammonium salt + base

salt + water + ammonia gas

Non metal oxide + base

salt + water

amphoteric metals and water + base

salt + hydrogen

base + dissolved amphoteric metal hydroxide

complex ions e.g Zn(OH)2 + 2OH- \= [Zn(OH)4] 2-

Amphoteric

- one that reacts with both acids and bases
- amphoteric metals include Al and Zn, which can react with bases to form hydrogen and also complex ions, e.g aluminate ion
- some insoluble metal oxides and hydroxide e.g Al(OH)3 can also dissolve in bases as well as in acids. Called amphoteric oxides and hydroxides.

Arrhenius Theory

- acid is a substance that produces H+ ions in solution
- alkali is a substance that produces OH- ions in solution
- strong acids are completely ionised in aqueous solution
- weak acids are partially ionised in aqueous solution
- strong bases are completely dissociated into ions in aqueous solution
- weak bases are bases which only a small proportion of base species react with water to form hydroxide ions
- Neutralisation: since H+ ions are responsible for properties of acids and OH- for properties of bases, neutralisation occurs when acids and bases react and these ions combine to form water

Issues with Arrhenius Theory

- restricts the acid base definition to solutions
- doesn't account for bases that don't have hydroxide ions

Bronsted Lowry Theory

- theory that supports the idea that acid base reaction is one that involves the transfer of protons from one species to another (H+ ion \= proton)
- acid is a substance that is a proton donor
- base is a substance that is a proton acceptor

Water in terms of Bronsted Lowry Theory

- if water donates a proton it acts as an acid and if it accepts a proton it acts as a base

Conjugate Acid-Base Pairs

in the bronsted Lowry theory:
- a base, after it has accepted a proton, has the potential to react as an acid (has a proton it can donate)
- an acid, which has donated a proton has the potential to react as a base (can accept a proton)
- every acid has a conjugate base: HX \= H+ ion + X- (conjugate base)
- stronger the acid, the weaker the conjugate base

Strongest Acid

HCl

Acid Base properties of Oxides

- both metal and non-metal elements can form oxides
- metal oxides are ionic e.g MgO, CuO
- non metal oxides are either covalent molecular or covalent network. e.g CO2, SiO2, H2O
- if an oxide is soluble in water, then pH can be tested to determine whether it is acidic or basic
- if an oxide is insoluble in water, then it must be reacted with an acid or base to determine whether it is acidic or basic

4 groups of Oxides

- Basic Oxides: display basic properties, not acidic
- Acidic Oxides: display acidic properties not basic
- Amphoteric Oxides: display both acidic and basic properties depending upon the reaction conditions
- Neutral Oxides: display neither acidic nor basic properties

Basic Oxides

- all basic oxides are metal oxides but not all metal oxides are basic oxides because some are amphoteric
- if the metal oxide is soluble in water, it reacts to form hydroxide ions
- basic oxides will also react with acids to form water and salt
- metal oxides are ionic and when dissolved in water, dissociate releasing positive metal ions
- oxide ion is a very powerful base that reacts with water to produce hydrogen ions

Acidic Oxides

- most acidic oxides are non-metal oxides. e.g SO3 + H2O \= H2SO4
- when mixed with solution containing OH-, acidic oxides produce water and polyatomic ion.

Ionisation of Water

- Water is weak electrolyte and to a very small extent undergoes self ionisation (H2O \= H+ ion + OH-, this is reversible)
- equilibrium strongly favours the reactants
- equilibrium constant (K) for ionisation of water is given by: K(w) \= [H+][OH-] \= 1.0x10^-14
- [H+] is concentration of hydrogen ions in mol per L
- [OH-] is concentration of hydroxide ions in Mol Per L
- K(w) is also called ionisation constant or dissociation constant for water
- in neutral solution, or pure water: [H+] \= [OH-] \= 1.0x10^-7
- when [H+] increases, [OH-] decreases

Strength vs Concentration

* strength of a solution is the determined by number of ions present and therefore the degree to which it is ionised or dissociated in water
* concentration of a solution is the amount of solute dissolved in a given volume of water

pH acidity scale

* pH of a solution indicates the acidity of a solution
* as \[H+\] increases, pH decreases
* Lower pH = greater acidity
* a neutral solution has a pH of 7
* change of 1 pH unit represents a 10 fold change in hydrogen ion concentration

Phases of matter

• Matter is described as anything that can take up space and has mass

• it exists in 3 physical phases: solids, liquids and gases

Kinetic Theory of Matter

Assumes:

• all matter is made up of tiny particles

• attraction between particles in condensed phases (solid and liquid) is strong and negligible in gases

• Particles of matter have kinetic energy due to their motion

• Average kinetic energy of particles defines their temperature

• particle collisions are elastic

Properties of Solid, Liquids and Gases

Using Kinetic Theory, factors that determine physical phase are:

• Inter-particle attraction

• pressure

• temperature

Understanding Gas Pressure

• kinetic theory can be used to explain why all gases exert pressure

• all gass particles move around inside a container they collide with the walls and in doing so exert a force on sides of the container

• this force, over a given area of surface represents pressure of gas

Effect of increasing amount of gas in a fixed volume

• more gas in a fixed container means more gas particles in the container. More particles means that there will be greater rate of collisions between gas particles

• Pressure increases with: increase with amount of gas in a fixed space, increase in temperature of the gas, decrease in gas volume

Pressure-Volume relationship

• the volume of a given mass of a gas is inversely related to the pressure exerted on it at a given temp and given number of molecules.

• P is inversely proportional to V, or P1V1 = P2V2

Pressure-Temperature relationship

• for a given mass and constant volume of an ideal gas, the pressure exerted on sides of its container is directly proportional to its absolute temperature

• P is proportional to T (in Kelvin), or P1/T1 = P2/T2

Volume-Temperature relationship

• for a given mass of an ideal gas at constant pressure, the volume is directly proportional to its absolute temperature in a closed system

• V is proportional to T (in Kelvin) or V1/T1 = V2/T2

Ideal Gas and Kinetic Theory

• kinetic theory of gases is used to explain physical behaviour of gases

• important to rmbr that this theory uses model of ideal gas

• Chemists apply kinetic theory to real gases under normal conditions

Properties of particles of a real gas

1. Occupy space and volume

2. have forces of attraction for another

3. Condense to form a liquid as its particles always have intermolecular attraction between molecules

Combined Gas Law

• (P1V1)/T1 = (P2V2)/T2

• PV = nRT is another formula

• where P1, V1 and T1 represent initial pressure, volume and temperature (K) and P2, V2 and T2 represent final pressure, volume and temeprature (K)

• R = 8.314J/K/mol where n = number of moles

• temperature must be in Kelvin and units for volume and pressure must be consistent.

100kpa

If the pressure of 1 atmosphere is 100kpa, then 300 kpa is the pressure of 3 atmospheres (approximately)

Temperature and Phase Change

• phase changes are physical processes and readily reversible

• phase changes are melting, sublimation, evaporation, boiling, condensation, freezing or deposition

• heating and/or reducing pressure can cause matter to change phase

• heating a substance causes increase in kinetic energy of substance

• at a particular temperature, the increased movement energy allows solid particles to overcome the attractive forces and move away from their fixed positions

• the solid melts and particles are able to move more freely.

Evaporation and Vapour Pressure

• evaporation occurs when particles in the liquid phase have a high enough kinetic energy to escape through the liquid surface to form a gas or vapour phase

• Because the more energetic particles are the ones that tend to escape, this will lower average kinetic energy of particles remaining in liquid

Vapour Pressure

• When a liquid is placed in a sealed container the amount of liquid will begin to decrease. This is because the more energetic particles in the liquid are escaping from surface of liquid to become a gas

• eventually, the amount of liquid present in container remains constant

• as the number of gas particles builds up in the container, the possibility of these gas particles colliding with the liquids surface and being attracted back into the liquid state (condensing) increases

• two processes are occuring simultaneously in the container

• the evaporation process for the liquid that occurs at a constant rate at a given temperature is taking place. Condensation of the vapour (gas) to reform some fo the liquid is also occurring.

• eventually, the rate of evaporation of liquid = rate of condensation of the vapour

• when no overall change is occurring in the volume of liquid in the flask the system is said to be in a state of dynamic equilibrium (two processes occurring at the same time).

Establishing A liquid-vapour equillibrium in a closed flask

• the gas particles that evaporate from the liquid exert pressure due to their collisions with the walls of the container

• This pressure of the gas in the equillibrium with the liquid in a closed container is called the equilibrium vapour pressure of the liquid, or more commonly, the vapour pressure of the liquid

• temperature of liquid also affects its vapour pressure

• as temperature rises, vapour pressure of liquid increases

• at higher temperatures a greater proportion of particles will have sufficient kinetic energy to escape from the liquid to produce the vapour

Boiling

• when water is boiled to 100C at 100 kPa (1 atmosphere), it starts to boil

• bubbles of steam can only form in the hot water when vapour pressure of the water = atmospheric pressure of surroundings

• The pressure of the vapour within the ‘boiling’ bubbles needs to be large enough to stop them from collapsing because of atmospheric pressure exerted at liquid’s surface

• at the BP of the substance, continuous vapourisation takes place throughout liquid

• Evaporation is where vapourisation takes place at the surface only

• variation in atmospheric pressure will result in change in BP

• lower the atmospheric pressure the lower the BP

Colligative Properties

presence of a dissolved solute:

• lowers vapour pressure of pure solution (vapour pressure reduction)

• raises the BP above that of pure solvent (BP elevation)

• lowers the freezing point below that of pure solvent (freezing point reduction)

Rate of Reaction

typically, the rate or speed of reactions increases with increasing temperature, concentration, pressure and sub-division.

Measuring Rate of Reaction

• Reaction Rate = amount of substance produced/time taken

• when observing rate of reaction graphically, the gradient gives an instantaneous rate of reaction at any point in time

Collision Theory

for a reaction to occur all the following conditions must be met:

• 1. particles must collide

• 2. collision energy must be equal to or greater then activation energy

• 3. reacting particles must collide with suitable orientation

• if reacting particles collide with sufficient energy and suitable orientation, they can form a transition state

• the bonds are breaking while new bonds are forming

• transition state exists for only a very short time.

Potential Energy Profile

• if colliding particles have sufficient kinetic energy, they can approach close enough to form a transition state (activated complex)

• transition state is highest potential energy state for a reaction

• activation energy = minimum collision energy required for a reaction

• stronger bonds means higher activation energy

Collision Energy and Temperatures

higher temperature means higher average collision energy of particles since temperature is a measure of average kinetic energy of particles in a substance

Factors affecting rate of reaction

Concentration, Gas Pressure, Temperature, State of Subdivision, Catalysts

Concentration (RoR)

• increase concentration = increased reaction rate

• higher concentration means there is a greater collision rate

• higher concentration of reacting particles colliding with correct orientation and sufficient activation energy

Gas Pressure (RoR)

• raising pressure (by reducing volume or adding more gas) crates a greater concentration of reacting gas molecules

• therefore, increased pressure = increased rate of collisions = increased rate of reaction

Temperature (RoR)

increased temperature means increased number of particles with sufficient energy as well as increased successful collisions and therefore increased reaction rate

State of Subdivision

• Heterogeneous reactions involve reactants in 2 separate phases (e.g solid/liquid)

• in these reactions particles can only collide at the surface boundary where separate phases make contact

• increased surface surface area = increased rate of collision = increased reaction rate

Catalysts

• catalysts: chemical substance that speeds up a reaction while remaining chemically unchanged at the end of the reaction

• catalysts increase reaction rate by providing a reaction pathway with a lower activation energy

• therefore, a greater percentage of collisions will have energy equal to or greater then activation energy