chem lecture test 1 ipt

Chem Lecture T1 15 JD TY N

Memory Work

What to memorise Definition Isotopes Have same proton number but different

nucleon number

Same chemical properties, but different physics properties

Atoms of an element

Ionisation energy First Ionisation energy is the amount of energy required to remove 1 mole of

electrons from 1 mole of gaseous atoms to form 1 mole of unipositive gaseous ions.

Second ionisation energy is the amount of

energy required to remove 1 mole of electrons from 1 mole of unipositive

gaseous ions to form 1 mole of gaseous ions with double positive charge.

Metallic Bonding Electrostatic force of attraction between metal cations and sea of delocalised

electrons Ionic Bonding Electrostatic force of attraction between

oppositely charged ions Covalent Bonds Electrostatic force of attraction between

nucleus of the two bonded atoms ahd shared pair of electrons

Explaining why dative bonds form Because one of the molecular has lone pairs and can act as an electron donor

while the other molecular has only 1/2/3 electron pairs, has an empty orbital in

valence shell and can accept the donated electrons to achieve octet.

Polarity Separation of Charge, unequal sharing of charge, dipole moments

Explaining solubility/interaction between ionic and covalent substances

Theory Substantiation

Covalent

Energy released from Solute solvent interaction greater than/smaller than

energy required to break solute solute and solvent solvent bonds

Covalent

Need to IDENTIFY the different types of

interaction in solute solute, solvent solvent, solute solvent before elaborating.


Then proceed to explain which bond is the stronger one and how it leads to the

solubility.

Ionic Compounds

Energy released from ion-dipole interactions >/< energy needed to break

hydrogen bonds of water and solute solute bonds, detaches ion from surface

Need to IDENTIFY the different types of

interaction in solute solute, solvent solvent, solute solvent before elaborating.

Then proceed to explain which bond is the

stronger one and how it leads to the solubility.

Sketching Important Trends

Trend Sketch+Explain Atomic Radius Across the Period

1. Atomic radius decreases down the period as - Nuclear charge increases due to increasing proton number - Shielding effect remains relatively constant - Effective Nuclear Charge increases - Electrostatic forces of attraction between ____ and _____

increases


- More energy required to overcome

2. Ionic radius generally decreases down the period, except for huge jump from cations to anions - Nuclear charge increases - Shielding effect remains relatively constant - Effective Nuclear Charge increases - Electrostatic forces of attraction between ____ and _____

increases - Sharp jump from cations to anions due to additional quantum

shell of electrons (3s and 3p)

3. Trends in successive I.E - Small increases when moving from p orbital to s as s orbitals at

higher energy level as they are CLOSER to the nucleus - Large jump when next electron is removed from an inner

quantum shell (look at the number of valence electrons)


4. Trends in IE across period - Nuclear charge increases across period - Shielding effect remains relatively constant - ENC increases - Electrostatic forces of attraction between ___ and ___

increases - More energy needed to remove outermost electron

Some anomalies - Small dip from Mg to Al as 3p subshell of al at higher energy

level than 3s of Mg, weaker attraction between nucleus and outermost electron, less energy required.

- Inter-electronic repulsion for paired electrons in 3p subshell of S, less energy required to remove the more unstable electrons.

Deducing strength of Bonds

Type of bond Explanation

Metallic -Compare number of valence electrons contributed PER METAL ATOM

(Stronger/weaker electrostatic forces of attraction)

-Charge, radius and charge density. Larger the charge and smaller the radius, ,the

stronger the bond. Ionic Lattice Energy (Charge of cation, anion,

radius of cation, anion)

Remember to break cations and anions up separately to compare and analyze.

Covalent Polarity, size and number of bonds (The smaller the atom, the more effective the orbital overlap, the closer to the nucleus.


The more polar, the shorter the bond length. The more number of bonds, the stronger the attraction between nucleus

and shared electrons.)

Explaining Chemical/Physical Properties

Type Properties Metals 1. High Melting and Boiling Points

- Exist as giant metallic lattice, strong electrostatic forces of attraction between metal cations and sea of delocalised electrons, more energy needed to break..

2. Good electrical conductivity - Sea of delocalised electrons to

act as mobile charge carriers 3. Good heat conductors

- Heat energy picked up by electrons to vibrate faster (kinetic energy)

4. Shiny surface - When photon strike surface,

osicillating electric field causes electrons on surface to oscillate. Photon bounces off without loss of momentum.

5. Hardness - Strong metallic bond

6. Malleable and ductile - Layers of ions can slide over one

another into new position. Sea of delocalised electrons reduces repulsions between cations, so do not break.

7. Alloy - Disrupt orderly arrangement of

metal atoms Ionic Compounds

1. High Melting and Boiling Points - Exist as giant metallic lattice,

strong electrostatic forces of attraction between metal cations and sea of delocalised electrons, more energy needed to break..


2. Good Electrical conductors in molten/aqueous state - Ions in solid state are held in a

fixed lattice, but when dissociated can act as mobile charge carriers

3. Hard, rigid but brittle - Hard force may cause ions of

like charges to become next to each other, causing repulsion and lattice to shatter.

Covalent 1. Low melting and boiling point for simple molecular structures and high for Giant covalent structures

- SMS molecules held by weak VDW forces

- GCS molecules held by strong covalent bonds

2. Most do not conduct electricity except for Graphite - As electrons are held in covalent

bonds/nucleus - As for graphite, each carbon is

only covalently bonded to three others, there is a spare electron to act as mobile charge carrier

Drawing Sigma/PI bonds Type of Bond

Drawing

Sigma Bond


Pi Bond

1s, 2s, 3s

Note: 2s orbital larger than 1s, 3s larger than 2s

Explaining difference between bond angles

Theory Explanation

1. Compare number of electron pairs first

2. Compare number of lone pairs if

electron pairs is the same

3. If all are the same, use polarity. A more polar molecular will have

slightly differing bond angles than a less polar one.

If electron pairs is different, then immediately use VSEPR to explain

Lone pair lone pair repulsion is the strongest, so it will push bond pairs closer. Remember the sequence= LP-LP repulsion> LP-BP Repulsion > BP-BP repulsion. Electrons in a more polar molecule will be attracted to __, leading to decrease in BP-BP repulsion (depends) and increase in LP-BP repulsion.

Determining Difference Between Melting/Boiling Points Key Possible ways of Answering Required


Remember that melting/boiling points are always due to intermolecular forces of attraction. Ionic compounds: Ions Metallic compounds: Ions and Delocalised electrons Covalent compounds: Molecules

Check for Differences in type of bonding

1. Comparing the structure (Ionic VS Covalent VS metallic compounds) - Giant ionic lattice, simple

molecular structure, giant molecular structure, giant metallic lattice

Check for Hydrogen Bonds

2. Hydrogen bonds are strong

- Requires FON with lone pair

- H bonded to FON

Confirm the Intermolecular Force

3. Polar or Non-Polar - Non-polar compounds have

dispersion forces, which are weaker than the permanent dipole permanent dipole interactions in polar compounds

Non Polar Compounds

Check for difference in Mr

4. If both molecules are non-polar, compare Mr - Larger Mr means

larger electron cloud, more polarisable, more formation of induced and instantaneous dipoles, stronger dispersion forces

Check for Branches

5. If both molecules are non-

polar, compare the surface area - A straight chain

compound has larger points of comtact that


dispersion forces can act. Therefore, stronger dispersion forces.

Polar Compounds

Difference in Polarity

6. Increase in polarity leads to shorter bond length - More polar means

shorter bond length, electron more attracted to nucleus, more energy needed to break bonds

chem lecture test 1 ipt

Documents