S3.2-3.3-3.1 ions, metallic, structure, alloys

Empirical formula vs molecular

empirical is the simpelest whole number ratio of atoms it contains (limited use as it doesnt tell us the actual number of atoms in the molecule)

molecular is the actual number of atoms of each element present (multiple of empirical formula)

If we have a ration between molar mass of empirical and molecular formula

then that is the same ratio for number of atoms in either formula, so if there were 100g/mol of empirical formula Hg4Fe3, and 300 g/mol of molecular formula, the molecular formula would be Hg12Fe9

Structural formula

A representation of the molecule showing how the atoms are bonded

Full structural formula- shows every bond and every atom

Condensed structural formula- omits bonds where they can be assumed, groups atoms together

Full structural formula would be the C-C with six H-s around it

condensed would be CH3CH3

Stereochemical formula

Shows the relative positions of atoms and groups around a central carbon in 3 dimensions

Dotted line goes back behind paper

solid line is in the plane of the paper

Thick triangle line come forward off the paper

Skeletal formula

shorthand representation of a structural formula, shows all bonds present in the molecule, except C-H bonds, omits C and H symbols

C4H10

Examples ethanol structural formula (C2H5OH), propanone (C3H6O), butanoic acid (C3H7COOH)

first sketch out how it should look normally, and then make skeletal formula, draw each C spike, and the extra things are another line coming off of it

Molecules containing a benzene ring

containing C6H6 we use the benzene ring to show these aromatic compounds

Catenation

Carbon is able to link to itself and form chains of bonded carbon atoms (why we have so many organic compounds)

Carbon atoms are able to form a variety of compounds containing single double or triple bonds

Functional groups

Atoms or groups of atoms that are present in organic compounds- responsible for a compound’s physical/chemical properties

compounds that contain the same functional group belong to the same class

ex class of carboxilic acids contain the same -COOH functional group

Common classes of organic compounds

Alkane-

suffix: -ane

ex ethane

general formula CnH2n+2

Alkene-

functional group- alkenyl

suffix-ene

ex ethene

CnH2n

Alkyne-

functional group- alkynyl

suffix-yne

ethyne

CnH2n-2

Alcohol-

funcitonal group- hydroxyl OH

suffix- anol

C2H5OH ethanol

General formula CnH2n+1OH

Ether

R-O-R’

functional group- alkoxy

suffix- oxyalkane

H3C-O-C2H5 methoxyethane

general formula- CnH2n+2+O

Aldehyde-

Functional group- carbonyl (aldehyde)

suffix- anal

ex of compound- C2H5CHO propanal

CnH2nO R-CHO

always at end

Ketone-

functional group- ketone

suffix- anone

ex. propanone CH3COCH3

general formula- CnH2nO R-CO-R’

always in middle

Carboxylic acid-

Carboxyl functional group

suffix- anoic acid

ex. of compound- C2H5COOH propanoic acid

general formula- CnH2n+1COOH

Ester-

Functional group- carboxyl (ester)

suffix-anoate

ex of compound- C2H5COOCH3 methylpropanoate

general formula - CnH2nO2

R-COO-R’

Amide-

Functional group- amido

suffix- anamide

ex. C2H5CONH2 propanimide

general formula CnH2n+1CONH2

Naming esters and ethers

Esters

Name by prefix being the compound single bonded to the O and the stem being the compound bonded double to the O

Ethers

2 alkyl chains linked by an oxygen atom, longer chain will be the stem and shorter chain is regarded as a substieunt and has prefix alkox, shorter chain-oxy-longer chain

naming amine and arene

Amine-

functional group- amino

suffix- anamine

C2H5NH2, ethylanamine

general formula-CnH2n+1NH2

Arene-

Functional group- phenyl

suffix benzene

C6H5CH3, methylbenzene, C6H5 is benzene, whatever added on becomes the prefix

Naming summary

Prefix- position, number, and name of substituents

Stem- Number of longest chain of C atoms

Suffix- Class of compound determined by functional group

saturated vs unsaturated compounds

Saturated contains only single bonds

Unsaturated contains double or triple bonds

Reaction pathway

Synthesize target compounds by organising several reactions in a sequence so the product of one is the reactant of the next, each step involves a functional group interconversion

ex. the reaction of 2 amino acids shows how functional groups on different molecules can react to give a new class of product

still has functional group on both sides so condensation can happen again to form a tripeptide and eventually a polypeptide

Homologous series

Organic compounds that are classified into families of compounds, members of each series have common features, for example the alkane homologous series

methane, ethane, propane

all members of a homologous series can be represented by the same general formula

some homologous series are characterized by the presence of a functional group for ex. alcohols shown below

methanol (CH3OH), ethanol (C2H5OH), propanol (C3H7OH)

homologous series properties

neighbouring members differ from each other by CH2 in alkane example, meaning molecular mass increases by a fixed mass

successive members have successively longer carbon chain causing a gradual trend in physical properties of members

the effect of the length of the carbon chain on the boiling point; gets higher as you get longer carbon chain length

same with melting point, density (stronger van der waal forces as induces more induced and instantaneous dipoles so more energy needed to break bonds present, stronger intermolecular forces )

Effect of polar vs non polar bonds/ intermolecular forces on physical properties

IUPAC nomenclature

Rules 1) identify the longest straight chain of carbon atoms

1- meth

2-eth

3-prop

4-but

5-pent

6-hex

2)identify the functional group- Determines the suffix to the name, replaces the “ane” in the parent alkane

class- refers to the type of compound

functional group- site of reactivity in the molecule

position of the funcitonal group is shown by a number inserted before the functional group ending

the chain is numbered starting at the end that will give the smallest number ot the functional group

ex. Propan-2-ol CH3CHOHCH3

but-1-ene- CH3CH2CHCH2

sometimes you dont need the number for ex. carboxilic acid is always at end of chain

3-identify side chains, side chains or functional groups are known as substieunts and are given the prefix check table

NH2 appears as a suffix and a prefix, when its the only functional group it will take the suffix (anamine), but if it is one of two or more it will be a prefix (amino)

if there are more than one substituent groups we use commas between the numbers of where they are and the prefixes di, tri, tetra, if there is more than one subtituent we place them in the name in alphabetical order ex. CCLH2CBrClCH3- 2,bromo,1,2,dichloropropane

Organic compounds with the same molecular formula can have different structures (not same name, same molecular formula)

ex. C4H10 can be butane or 2-methylpropane, because of this they will have different properties

Same molecule, different arrangements of atoms are structural isomers

each isomer is a distinct compound with unique physical and chemical properties

number of isomers for a molecular formula increases with the size of the molecule, ex pentane has pentane, 2methylbutane, 2,2,dimethylpropane

the more branching, the lower the boiling point so pentane has the highest, 2methylbutane middle and 2,2dimethylepropane has lowest

Alkenes can have structural isomers

contain atleast 1 carbon carbon double bond, different type of isomer, not due to branching like alkanes but due to position of the double bond (functional group)

ex C6H12 hexene

also applicable for alkynes with atleast one triple bond and its possible positions

structural isomers can have functional groups attatched but they must be in the same place

positional isomers

have the same functional group but at different positions 1bromopropane vs 2bromopropane

metal atoms

lose electrons to form positive ions (cations)

non metal atoms

gain electrons to form negative ions

atomic number of an element is defined in, valence electrons vs protons

number of protons which does not change in the chemical reaction, valence electrons are the furthest from the electrostatic attraction of the nucleus so more open to external influences

Outer electrons of metal atoms experience a smaller effective nuclear charge than the outer electrons of non metals

nuclear charge given by atomic number which increases across a period

valence electrons which determine chemical properties do not experience full attraction of the charge as they are shielded from nucleus/repelled by inner electrons, presence of innershielding reduces attraction of the nucleus for outer electrons and effective charge experienced by valence electrons is less than full nuclear charge

across period ENC increases as same shielding but increase in proton number

down a group increase in nuclear charge from increase proton number is offset by increase in inner shielding remains the same

Metal atoms form positive ions as they (cations)

have low ionization energies (little energy needed to remove outermost electron)

first ionization energies increase across a period due to increase in ENC

increased ENC, harder to remove electron

more likely to lose an electron and form a positive ion if it is a metal on left of periodic table, low ionization energy

ionization energy decrease down a group metals on bottom left have greatest tendancy to lose electrons and form ionic compounds

Non metal atoms form negative ions as they have (anions)

a high effective nuclear charge

electron transfer will be more likely to occur if non metals attract the transferred electrons more strongly

non metal top right highest ENC (most atomic number and least inner shielding layers) and smalledt atomic radii

once electron transfer is complete

electrostatic attraction between oppositely charged ions pull the ions together

most vigorous reactions occur

between elements furthest apart in the periodic table

caesium is the most reactive alkali metal, bottom group 1, lowest ionization energy

flourine is the most reactive non metal, smallest atomic radius, attracts transferred electrons the strongest

increased force of attraction ionic

greater charge, more electrons transferred

halogens will attract electrons the strongest as they have only one vacancy left to fill their outer shells

formation of Si4- ion not feasible because

the addition of electrons becomes more difficult with increasing negative charge of the ion due to increased electron electron repulsion

High first ionization energies

(lot of energy needed to remove electron), noble gases, nucleus holds tightly to outer electrons, not available for chemical activity, complete outershell, added electrons would occupy an empty outer energy level, ENC of zero

Transition metals form

ions of different charge, iron (Fe2+ or Fe3+) or copper (Cu+ or Cu2+)

the iron and copper ions of different charge form compounds with different properties like color

Naming for transition metals using oxidation number

only necessary when an element has more than one oxidation state, ex Na2O only shows one in compounds

Formula compound- FeO, oxidation state Fe2+, name using oxidation number Iron (II)oxide

Cu2O, Cu+, Copper(I)oxide

CuO, Cu2+, copper (II)oxide

MnO2, Mn4+, manganese (IV)oxide

ionic compound is generally nuetral, how to balance negative and positive charge

look at charge of both, and see what you need to multiply one or both by to get the lowest same whole number ratio

Polyatomic ions

made up of more than one atom together have lost or gained an electron, found in salts formed from common acids

NO3- HNO3 acid

NO3- Nitrate

SO42- sulfate

PO43- phosphate

OH- hydroxide

CO32- carbonate

HCO3- hydrogen carbonate

NH4+ ammonium

CHC3OO- ethanoate

bonds holding polyatomic ions together are covalent bonds, but once they join together, the bonds that joins two polyatomic ions are ionic bonds

Ionic compounds have a lattice structure, ionic bond, geometry

once ions are formed by electron transfer they are pulled together by electrostatic attraction which is the ionic bond

many cations and anions arrange themselves in a 3D lattice structure held together by ionic bonds between oppositely charge ions

geometry varies for each compound depends on sizes of ions, involves a fixed arrangement of ions based on a repeating unit cell

consists of a very large number of ions, can grow indefinitely

ionic compounds do not

have a fixed number of ions so their formulas are ratios of ions present , empirical formula known as the formula unit

formation of the ion energetically feasible

energy output when ion is formed (neg enpalthy) is offset by energy required to form the ion (pos enpalthy)

Lattice enpalthy

the measure of strength of the ionic bond in different compounds influenced by ion radius and charge (amount of energy to seperate one mole of solid ionic compound into consituent gaseous ions)

enpalthy change can be calculated by

the ionic model which assumes that the crystal made up of spherical ions which only interact by electrostatic forces

increase in ionic charge increases the attraction between ions and increases lattice enpalthy

increase in ionic radius of one of the ions decreases attraction between ions and decreases lattice enpalthy

lattice structure results in

ionic compounds being crystalline solids at room temperature, high meleting and boiling points (large amount of energy needed to seperate the ions in the lattice), and low volatility

ionic solubility

high solubility in water as water is polar and the partial positive charges on H attract negative ions and the partial negative charges on O attract positive ions

at the contact surface of a crystal, attraction of ions to partial charges in water molecule pulls ions away from their lattice positions as those ions are now seperate they are surrounded by water molecules and are hydrated and dissolved

if liquid is non polar there is no attraction between liquid molecules and ions as there are no partial charges so ions remain in lattice and solid is insoluble

exceptions like calcium carbonate not soluble in water

solvents other than water, solute ions are said to be

solvated

ionic compounds electricity conductors

not able to conduct electricity in their solid state as the ions are now fixed within their lattice but they can conduct in the liquid state as the ions are mobile (can carry pos or neg charges from one area to another)

ionic compounds are brittle

the crystal structure shatters with shear force as ions within the lattice are displaced resulting in ions of the same charge being positioned alongside each other and repulsive forces between the pos-pos and neg-neg ion cause the lattice to split

when two elements react to form an ionic compound they show different tendancies to lose or gain electrons

metals on the bottom left lose electrons most easily, non metals on top right gain electrons most easily, caesium and flourine are the compound with the most ionic character

down a group, in first group most reactive metal, electrons furthest away from nucleus so most reactive and most tendancy to form positive ion, electron can be removed easiest

across a period, increasing tendancy to form negative ions, most reactive non metal group 17 period 1, electrons are closest and ionization energy is the highest and electron affinity is the highest

period 3 chloride compounds

less ionic across a period and more covalent, meaning lower melting point as turns more covalent (molecular covalent), and electrical conductivity in molten state gets lower (ions can carry pos or neg charge and move in ionic compound, atoms are nuetral and no charged particles free to move in molecular covalent), ionic bonds stronger than weak intermolecular forces

A covalent bond forms by

Atoms sharing electrons

Atoms of two non metal reacting together

Both seeking to gain valence electrons to achieve the stable electron structure of a noble gas

share electron pair to do this

shared pair of electrons is concentrated in the region between the two nuclei and is attracted to them both

electrostatic attraction between the shared pair of electrons and positively charged nuclei holds atoms together- covalent bond

why the atoms are held at a fixed distance apart in covalent bond

system containing the two atoms is stabilized when the forces of attraction between the nuclei and shared electrons are balanced by the forces of repulsion of the two nuclei

covalent bonds form at the point

of lowest energy as two atoms approach each other

octet rule

8 electrons for a full outer shell to predict stable arrangements in covalent bond

ex. cl needs to gain 1 electron because group 17 have 7 valence electrons, so both will share 1 electron

ability of two identical atoms to form a covalent bond

is due to the similair strength with which they attract valence electrons

atoms of group 18

have a complete octet without needing to share electrons, low levels of reactivity, do not form covalent bonds (except helium)

Incomplete octet

Be beryllium and B boron forms stable molecules in which central atoms have less than 8 valency electrons (small atoms)- limits number of atoms that can get close enough to share required number of electrons for complete octet

Be has 4

B has 6

Lewis structure

Describes the structure of covalent molecules

1- calculate total number of valence electrons in the molecule

2- look how many time each one has to bond (how many extra electrons it needs to gain)

3- all other electrons are lone pairs

4-check that all atoms have a full valency shell and that the total number of valency electrons is correct

if there are ions, if it is negative, you add an electron if positive you remove one, see where it fits best so that there are lone PAIRS of electrons and not an uneven number, and that there are still full valency shells and the right number of electrons

diatomic molecules

normally bond with double or triple bonds which are hard to break making the molecule very stable (not reactive)

multiple bonds have a greater number of shared electrons and a stronger force of electrostatic attraction to bonded nuclei, greater pulling power between atoms brings them closer together resulting in shorter and stronger bonds than single bonds

3x bonds is most 1x bonds is least strong

1x bonds is greatest length 3x bonds is least

oxygen, nitrogen, hydrogen, flourine, chlorine, bromine, iodine

Covalent bond characterized by

bond length- distance between 2 bonded nuclei

bond strength- measure of energy required to break the bond (enpalthy)

atomic radius increases down a group

more shells, then the bond length increases

bond enpalthy decreases because the shared electron pair is further from the pull of the nuclei in the larger molecule, so bond is weaker and takes less energy to break

Coordination bond, ex NH4+, H3O+, CO

A coordination bond is a covalent bond in which both shared electrons come from same atom

an arrow on the head of the bond is used to show coordination bond

once coordination bonds are formed, they are no different from other covalent bonds

shape of a molecule is determined by

repulsion between electron domains around central atoms

based on VESPR theory-

Because electron pairs in the same valence shell carry the same charge they will repel each other and spread out as much as possible

the repulsion applies to electron domains, which can be single, double, triple bonding electron pairs, or non bonding pairs of electrons

the total number of electron domains around the central atom determines the arrangement of the electron domains

the shape of the molecule is determined by the angles between the bonded atoms

non bonding pairs and multiple bonds cause slightly more repulsion than a bonding pair because a non bonding pair has a higher concentration of charge than bonding pairs, as they are not shared between 2 atoms and multiple bonds have a higher concentration of charge as they contain 2 or 3 pairs of electrons

electron domain

all electron locations in the valence shell, whether they are occupied by the non bonding pairs, single, double, triple bonded pairs

total number of electron domains and this can be determined from Lewis formula determines shape

repulsion force

single bond<double or triple bond<lone pairs

molecules with lone pairs or multiple bonds on central atom have some distortions in their structure that reduce the angle between the bonded atoms

electron geometry vs molecular geometry

electron geometry is based on the lone pairs and the bonding pairs, lone pairs is only on central atoms

molecular geometry is only the bonding pairs

2 electron pairs

Positioned 180 degrees from each other, linear shape

if there are two bonding pairs and two lone pair, molecular geometry is v shaped and electron geometry is tetrahedral (104.5 degrees if there is two lone pairs )

if there are two bonding pairs and one lone pair, electron domain is triangular planar, molecular geometry is v shaped (117 degrees)

3/4 electron domains

positioned 120 degrees from each other, triangular planar shape as electron domain geometry

when you have a double bond, bond angles will be slightly distorted due to increased repulsion from double bond, so less bond angles

if there are three bonding pairs and one lone pair, electron domain is tetrahedral and molecular geometry would be triangular pyrimidal, 107 degrees

tetrahedral is 109.5 degrees

4 electron domains

positioned 109.5 degrees from each other, tetrahedral electron domain geometry, same for molecular geometry if all 4 are bonding

Polar bonds occur from

If electron pairs spend more time with one atom than the other, they are not equally shared occuring when there is a difference in electronegativity of bonded atoms

electronegativity- ability of an atom to attract electrons in a covalent bond

as the more electronegative atom exerts a greater pulling power on the shared electrons, it gains more posession and electron distribution is uneven- polar bond

Bond dipole

This type of bond has two partially seperated opposite electric charge

The more electronegative atom with greater share of electrons is partially negative, less electronegative becomes partially positive

extent of polarity in a covalent bond varies depending on how big a difference between electronegativity values between bonded atoms

electronegativity periodic trends

increases across a period, more protons, same inner shielding, more ENC

Increases up a group, less inner shielding, more ENC

flourine is most electronegative atom, will have greatest electron density when covalently bonded to another element

only bonds that are non polar

are bonds between the same atoms, such as F2, O2, H2, difference in electronegativity is zero so it is pure covalent

C-H very low polarity, determines properties of organic compounds

polar vs covalent vs ionic

the more polar the bond the greater the seperation of charges, more like an ionic compound, so polar bonds act like an intermediate between pure covalent and ionic bonds which is why we see an overlap in properties of ionic and covalent substances

more than 1.7 ionic

more than 0.4 POLAR

Molecular polarity depends on

the polar bonds it contains

the way which such polar bonds are orientated with respect to each other; molecular geometry

vs. bond polarity depends on charge seperation between its two bonded atoms

if bonds are of equal polarity and are arranged symmetrically, their charge seperations will oppose/cancel each other out, the molecule will be non polar but contain polar bonds

ex. CO2

notation for a dipole that results from the pull of electrons in bond towards more electronegative atom is the arrow with a line

If either the molecule contains bonds of different polarities or its bonds are not symettrically arranged

dipoles will not cancel out and molecule will be polar and will have a net dipole

find which molecules are polar or non polar by

first draw lewis structure

find electronegativities of each molecule in data booklet

find differences in electronegativities and write them next to the bond

largest difference will pull most in that direction and have a partial negative charge

most covalent structures exist as

discrete molecules with a finite number of atoms

covalent network structure

some substances have a crystalline lattice structure in which the atoms are linked together by covalent bonds , single molecule with a regular repeating pattern of covalent bonds, has no finite size, have different properties than other smaller covalent molecules

Allotropes

have different bonding and structural patterns of the same element in the same physical state- different chemical and physical properties ex. molecular oxygen and ozone both exist as gases are examples of allotropes of oxygen

Carbon’s allotropes- diamond

structure- Each C atom is covalently bonded to 4 others, tetrahedral, repetitive pattern, 109.5 degrees bond angles

non conductor of electricity- all electrons bonded so non mobile

thermal conductivity- very efficient better than metals (strong covalent bonding)

appearance- transparent

physical/chemical- hardest natural substance, brittle, very high melting point

uses- tools for cutting glass or jewelery

Graphite

Structure- each C atom covalently bonded to 3 others, hexagons in parallel layers, bond angles in 120 degrees, valence electrons remaining delocalized and can move freely across layers held together by weak london dispersion forces so they can slide over each other

electrical conductivity good, contains one delocalised electron pair per atom giving mobility

thermal conductivity- not good unless heat can be forced to conduct in a parallel direction to crystal layers

appearance- grey crystalline solid

physical/chemical- soft and slippery due to slippage of layers over each other, brittle, high melting point, stable

uses in pencils

functional group isomers

molecules with the same molecular formula, but different functional groups

graphene

structure- each C atom is covalently bonded to 3 others, bonding angles of 120 degrees, 2D, remaining electron on each carbon is delocalized

electrical conductivity- very good, one delocalized electron per atom gives electrical mobility

thermal conductivity- best thermal conductivity

appearance- almost transparent

physical/chemical- thickeness of one atom, thinnest material, strongest, flexible, high melting point

uses— transmission electron microscopy

Fullerene

structure is carbon bonded to 3 others, closed spherical cage of 60 carbon atoms

electrical conductivity- poor conductors, little movement of electrons between molecules despite delocalised electrons

thermal conductivity- very low

appearance- black powder

physical/chemical- light strong low melting point

uses lubricants for medical purposes

graphite occurs naturally

single seperated layer is graphene, rolled up layer of graphene is a nanotube, closed cage is fullerene

development of structures like nanotubes, nanobuds, graphene is part of nanotechnology science through atomic scale and manipulation of matter

silicon

atoms have 4 valence shell electrons

in elemental state each is covalently bonded to four others in tetrahedral, giant lattice structure

silicon dioxide

forms a giant covalent structure based on tetrahedral arrangement

each Si atom covalently bonded to 4 O atoms and each O atom to 2 Si atom

strong,insoluble,high melting point, non conductor of electricity due to strong tetrahedral position, glass, sand

silicon commonly bonds with O and not with itself like carbon, which are stronger than si-si bonds as c-c bonds have a smaller atomic radius, greater electrostatic attraction

in a covalent lattice structure

covalent bonds hold atoms together within the molecule but intermolecular forces exist between the molecules depending on size and polarity of the molecules

strength of intermolecular forces determine physical properties of a substance (volatility, solubility, conductivity)

London dispersion forces

non polar molecules like Cl2 have no permanent seperation of charge within their bonds because shared electrons are pulled equally in both directions, no permanent dipole

electrons behave as mobile clouds of negative charge, the density of cloud may be greater over one atom than over another at any moment, when this happens an instantaneous dipole gives some seperation of charge on one atom, only lasts an instant, and may influence electron distribution in bond of a neighbouring molecule, induced dipole (makes the nearby side positive because the negatives repel away from each other)

weak london dispersion forces occur between opposite ends of two temporary dipoles in the molecules, weakest intermolecular force, strenght increases with increasing molecular size as you have a greater number of electrons which increases the probability and magnitude of instantaneous dipole formation

only forces between non polar molecules, low melting and boiling points, little energy required to overcome and seperate molecules

strength increases with size so does melting and boiling point, responsible for the fact that non polar substances can condense to form liquids at low temperatures

can also be components of the forces between polar molecules but get overlooked because other forces are stronger

Dipole dipole attraction

Polar molecules like HCl have a permanent seperation of charge (mismatching electronegativities)

one end is partial negative if it is electron sufficient and one end is partial positive if it is electron sufficient so there is a permanent dipole resulting in opposite charges on neighbouring molecules attracting each other

strength of this force will vary depending on distance and relative orientation of dipoles always stronger than london dispersion due to permanent rather than instantaneous dipole, melting and boiling points of polar compounds higher than those of non polar substances of same mass, leads to solubility of polar solutes in polar solvents

dipole induced dipole attraction

mixture between polar and non polar molecules, the permanent dipole of a polar molecule can cause a temporary seperation of charge on a non polar molecule- force is dipole induced dipole attraction (suddenly because there are so many electrons on one side of a neighbouring molecule, the other side’s electrons will be repelled away and positive charge will be attracted to neighbouring molecule’s negative side)

acts in addition to london dispersion forces between non polar molecules and dipole dipole forces between polar molecules

Van der Waals forces

Includes all three forces, refers to all forces between molecules that do not involve electrostatic attractions between ions

van der waals forces in some cases can occur within a molecule if different groups can position themselves appropriately

Hydrogen bonding, why ice floats on water, why water is liquid at room temp

Molecule containing hydrogen covalently bonded to a very electronegative atom

strongest intermolecular force

particular case of dipole dipole, large electronegativity difference between hydrogen bonded oxygen for ex which causes electron pair to be pulled towards oxygen, becuase of its small size and having no other electrons to shield its nucleus, hydrogen with a partial positive charge exerts a strong attractive force on negatively charges lone pair in the electronegative atom on the neighbouring molecule with a partial negative charge

higher melting and boiling points than other forces with same mass

if not for hydorgen bonding, water should be a gas at room temperature due to its low mass, but because of two hydrogen atoms and 2 lone pairs, each water molecule can for four hydrogen bonds with neighbouring molecules but liquid will contain less than four

ice contains four and is maximally hydrogen bonded resulting in tetrahedral arrangement holding molecules at a fixed distance apart and less dense than liquid so that way it floats on water

density change means water expands on freezing

hydrogen bonds weaker than

covalent and ionic bonds but stronger than all other intermolecular forces

intermolecular and intramolecular forces both are categories of electrostatic attraction, but significant difference in relative strengths

Intermolecular forces table

melting and boiling points, what are they, covalent vs ionic, covalent giant vs molecular, molecular size and polarity, hydrogen bonding

changing states by melting and boiling involves seperating particles by overcoming the forces between them

the stronger the interparticle forces, the more energy that will be required to break the forces and higher melting or boiling points

covalent substances have lower melting/boiling points than ionic, forces to overcome to seperate molecules are the weak intermolecular forces, easier to break than the electrostatic in ionic lattice

strength of intermolecular forces increases with increasing molecular size (more electrons, more magnitude of intermolecular forces) , also with increase in extent of polarity due to more dipole forces

simple covalent molecules have a small and fixed number of atoms, easier to break intermolecular forces in these than giant molecular structures have stronger covalent bonds (repeating lattice is harder to break), can be as strong as ionic compounds

Hydroxyl functional group leads to formation of hydrogen bonding in acohols

volatility

Tendancy of a substance to vaporize (liquid/solid to vapour)

substance with strong intermolecular forces will have a lower tendancy and vice versa

C3H7OH is less volatile than C4H10 due to strong hydrogen bonding

ionic compounds- low, giant covalent- low polar- high non polar-highest

solubility

non polar substances dissolve in non polar solvents by london dispersion forces between solute and solvent

all halogens which are diatomic and non polar, readily soluble in non polar (parrafin oil solvent)

Polar covalent compounds are soluble in water (polar solvent) interact through dipole interactions and hydrogen bonding ex. HCl in water

solubility of polar compounds is reduced in larger molecules where polar bond is a small part of total structure, non polar parts reduce its solubility ex. C2H5OH is polar, and C7H15OH is non polar

polar susbtances have low solubility in non polar solvents, remain held to each other by dipole dipole attractions and dont interact well with solvent

giant molecular structures insoluble in all solvents as there is too much energy required to break the strong covalent bonds

ionic compounds are polar extremely due to the transfer of electrons so they are soluble in water, non soluble in non polar susbtances

electrical conductivity

covalent compounds do not contain ions and so are unable to conduct electricity, no mobile ions , both non polar covalent and giant covalent

some polar covalent compounds in conditions where they can ionise, will conduct like Hcl in water that dissolves into its H plus ions in water

ionic compounds conduct when molten or dissolved in water

Chromatography

Seperate and identify components of a mixture

components have different affinities for two phases and are seperates as mobile phase moves through stationary phase

levels of solubility of each component dependent on its intermolecular forces

paper containing 10% water is stationary phase

solvent is the mobile phase as it rises up the paper by capillary action dissolving the components of the mixture to different extents carrying them up at different rates

trial multiple solvents to identify which one seperates the components best

polar compounds will stick to stationary phase and not move up with mobile phase when solvent is non polar

non polar substances will move up with non polar solvent mobile phase