Week 3- Atoms, Ions and molecules (Periodic Behaviour, Ionisation, Oxidation States, Covalent Bonding)

Question 1

Octet Rule

Answer 1

“happiness is a filled shell”

-All elements strive become a noble gas. Filling outer shell (by gaining or losing)

Question 2

What does a filled and unfilled valance shell mean?

Answer 2

Filled - no bond formations
Unfilled - reactivity

Question 3

The Lewis Model of Bonding:

Answer 3

Model uses valence electrons to derive bonding structures.
key is each atom involved in bonding, should attain configuration of noble gas

Question 4

Lewis Atomic Structure:

Answer 4

Atoms indicated with their normal symbols.
Valence electrons represented by dots

N.B Hand’s rule must also be used (i.e most stable configuration, one with most unpaired electrons.

Question 5

Stable Electron Configuration:
Two consequences

Answer 5

• Atoms will try get into stable configuration ( chemical bonding)

-Different elements may get to stable configuration in different ways (electronegativity)

Question 6

Driving force for bonds

Answer 6

-Chemical bonds make atoms increase in stability then if non-bonded

-Bond formation equals change in electrons, through electron transfer or sharing

Question 7

Ions

Answer 7

atom or group atoms, has charge

Question 8

Cations

Answer 8

positively charged atoms

Question 9

Anions

Answer 9

negatively charged atoms
(non metals)

Question 10

Rules to predict charge:

Answer 10

Cations - charge = main group no.

Anions - charge = (8 - main group no.)

N.B if ion has ide in name = anion

Question 11

Atomic size

Answer 11

Measure atomic radius
Distance between 2 nucleuses (2r) equals the atomic size

N.B only applicable if atoms must be identical (e.g a molecule)

Question 12

What is the general trend for Atomic size in periodic table

Answer 12

Trend: Increases top to bottom, Decreases left to right

Question 13

What is the group trend for Atomic size in periodic table
(explain)

Answer 13

Atomic radius increases as atomic no. increases.

Shielding effect increases, due to no. energy levels.
Shielding effect > nucleus charge attraction therefore atomic size increases

Question 14

What is the periodic trend for Atomic size in periodic table
(explain)

Answer 14

Left to right, no. electrons added to same principal energy level increases.
Shielding effect constant, but there’s increase in nuclear charge, pulls highest occupied energy level closer, atomic size decreases.

Question 15

Electronegativity

Answer 15

measure, tendency atom attract a bonding pair of electrons

Question 16

Electronegativity general trend in periodic table

Answer 16

Increases bottom to top
Increases left to right

Question 17

Ionisation Energy

Answer 17

energy required remove electron from atom (only measured when element in its gaseous state)

Question 18

First Ionisation Energy

Answer 18

energy required remove first electron from an atom

second ionisation energy etc

Question 19

Can we predict what ions an element will form?

Answer 19

yes

Question 20

General trend in Ionisation energy, in periodic table

Answer 20

Increases bottom to top
Increases left to right

Question 21

Lewis Diagrams (video)

Answer 21

Shared electrons are counted as owned by both atoms.

1 line = single bond = 2 electrons
2 lines = double bond = 4 electrons
3 lines = triple bond = 6 electrons
2 dots are called a lone pair of electrons

Question 22

Steps for drawing a Lewis Diagram (video): (5)

Answer 22

1) Count all the VALENCE electrons
2) Determine the central atom (the element there is only one of)
3) Draw single bonds to the central atom
4) Put all remaining valence electrons on atoms as lone pairs (dots)
5) Turn lone pairs into double or triple bonds to give every atom an octet (or deut)

Question 23

Electronegativity Definition: (video)
What makes an element more electronegative?

Answer 23

How well an atom can attract shared electrons

Dependant on it’s no. of protons, pull electrons closer as it has a higher nuclear attraction

Question 24

What’s the difference between nuclear charge and effective nuclear charge?

Answer 24

The nuclear charge is a measure of how positive
the nucleus is. The higher the atomic number of an
element, the more protons it has in its nucleus, and
hence the higher the nuclear charge.

The effective nuclear charge is the net positive
charge experienced by the outermost electron from
the nucleus. It is less than the nuclear charge,
because it takes into account the repulsion from
inner electrons which provide a shielding effect.

As you move from element to element across the Periodic
Table, you are adding extra protons to the nucleus, and
extra electrons around the nucleus.

Question 25

Ionic size

Answer 25

• During reactions between metals and non-metals, metal atoms lose electrons
  non-metal atoms gain electrons.
• Tansfer electrons effect on size ions that
  form.
• Cations smaller than the atoms from which they form.
• Anions always larger

Question 26

Trends in ionic size - summary

Answer 26

Size increases going down
Size decreases going right for cations and anions

Question 27

Periodic trends - summary

Answer 27

Atomic size, ionisation energy, ionic size, and electronegativity are trends that vary across
periods and groups of the periodic table.

• These trends can be explained by variations in atomic structure.
• The increase in nuclear charge within groups and across periods explains many trends.
• Within groups, an increase in electron shielding has a significant effect on these trends.
• In general, atomic size increases from top to bottom within a group and decreases from left to right across a
  period.
• In general, electronegativity values decrease from top to bottom within a group. For representative
  elements, the values tend to increase from left to right across a period.
• Positive and negative ions form when electrons are transferred between atoms.
• First ionisation energy tends to decrease from top to bottom within a group and decrease from left to right
  across a period.
• Ionic size tends to increase from top to bottom within a group. Generally, the size of cations and anions
  decreases from left to right across a period.

Question 28

Properties of ionic compounds (6)

Answer 28

• Hard, rigid solids at room temperature
• High melting point
• Dissolve in polar solvents (if soluble)
• Solutions conduct electricity
• Melts conduct electricity
• Closely packed dense structures

Question 29

Units can be…

Answer 29

atoms or ions

Question 30

Metallic bonding
Metallic:

Answer 30

All atoms form positive ions by getting rid of electrons,
which in turn are delocalised to form a “sea of electrons”.

Question 31

Properties of metals (6)

Answer 31

• High melting points
• Good conductors of electricity
• Good conductors of heat
• High density
• Malleable
• Ductile

Question 32

Electron shielding:

Answer 32

electron shielding refers to the blocking of valence shell electron attraction
by the nucleus, due to the presence of inner-shell electrons. Electrons in an s orbital can shield p
electrons at the same energy level because of the spherical shape of the s orbital.

Question 33

When drawing dot structures for compounds what element goes in the centre?

Answer 33

The less electronegative

Question 34

Make sure the formal charges for an element when drawing the lewis diagram is preferably equal to …

Answer 34

zero

Question 35

Sometimes formal charge will effect …

Answer 35

the final strucutre of your molecule or ion

Question 36

Formal charge out plays what rule

Answer 36

octet rule

Question 37

Steps for predicting and drawing structure of a molecule / ion using VSEPR (4 steps):

Answer 37

1) Draw a dot strucutre to show the valence electrons

2) Count the no. of electron clouds surrounding the central atom

3) Predict the geometry of the electron clouds around the central atom (like charges and so clouds repel as much as possible)

4) Ignore any lone pairs and predict the geometry of the molecule / ion,

Question 38

How does metallic bonding work?

Answer 38

Metallic: All atoms form positive ions by getting rid of electrons, which in turn are delocalised to form a “sea of electrons”.

Question 39

What does binding strength in metallic bonding depend on?

Answer 39

on # of delocalised electrons

Question 40

Metals can be pure metals (repeating organised structure of one element), but they may also be compounds (alloys): what are 2 examples

Answer 40

Interstitial Alloy (one smaller type of metal, sitting in between lattice of bigger metal)
Substitutional Alloy (when two types metal similar in size).

Examples of alloys are steel, brass and bronze

Question 41

What does the type of covalent bond depend on?

Answer 41

the # of shared electrons and their source

Question 42

What are the different types of covalent bonds:

Answer 42

single bond, One pair of shared electrons
double bond, Two pairs of shared electrons
triple bond, Three pairs of shared electrons
Dative bond, Both electrons contributed by one atom.

In dative covalent bonding (also known as coordinate bonding) the atom that supplies the electrons is known as the donor, while the atom that receives the electrons is known as the acceptor. An arrow is usually used rather than a line to indicate the acceptor and donor species

Question 43

What is bond enthalpy

Answer 43

represents the energy needed to break a specific bond (could be single, double, etc)

SI units of energy is the Joule (J)
or Joules per mole

Question 44

Shorthand electron configurations

Answer 44

Use noble gas symbols

We can do so because, we are only interested in the valence shell electrons for structure prediction as the core electrons will not be involved (they satisfy the octet rule).

Question 45

Criteria for covalent bond formation: (2)

Answer 45

(1) The energy of the valence electron on A should be similar to that on B. If there is a significant difference then electron transfer could occur resulting in an ionic bond A+ B-

(2) For the electron pair to be strongly attracted to the positively charged nuclei, it is preferable if the
electron pair is localised between the nuclei and that they are spin paired (i.e. ms = +1/2 and –1/2).

(3) For electrons to occupy the same region of space between the nuclei the orbitals of A and B must overlap for a bond to be made.

(4) The covalent molecule will seek the lowest possible energy. This means that:
* the maximum number of bonds will form
* the strongest possible bonds will form and
* the geometrical arrangement of these bonds will be such that electron pair repulsion is minimised.
(more on this later)

Question 46

The Octet Rule and Duplet Rule:

Answer 46

An atom obeys the octet rule when it gains, loses or shares electrons to give an outer shell containing eight electrons with the configuration ns^2np^6, i.e. the nearest noble gas configuration. Note when H bonds covalently it obeys the duplet rule, i.e. it achieves the [He] electronic configuration of 1s^2

Most second period elements (Li to F) exist in compounds in which each atom is surrounded by eight electrons.

Question 47

What does bond order equal?

Answer 47

Bond order = Number of bonding pairs of electrons

Question 48

Heterodiatomic molecules, example

Answer 48

Carbon requires four electrons to achieve an octet whilst oxygen requires two. The bond order in CO is also three.

Question 49

Charged species:
If a species is charged the number of valence electrons available for bonding is simply adjusted by:

Answer 49

· removing one electron for every unit of positive charge
·
adding one electron for every unit of negative charge

Question 50

Steps to drawing a Lewis structure:

Answer 50

Remember the Octet rule – In many molecules each atom shares electrons with
neighbouring atoms to achieve 8 valence electrons (an octet).

a) identify the total number of valence electrons present in each contributingAtom

b) if the molecule is charged subtract or add extra electron(s) to account for the positive
or negative charge respectively

c) distribute electrons in the structure so that there is an electron pair forming a single
bond between each pair of atoms that are bonded together. The extra electrons are then
added as lone pairs or form multiple bonds until each atom has an octet

Question 51

What does VSEPR Theory stand for, what is it used for?

Answer 51

VSEPR – Valence Shell Electron Pair Repulsion.

Predicting Molecular Shape from Lewis dot molecular structures

Question 52

Definition of Resonance

Answer 52

when several alternative LOW ENERGY Lewis structures are available the bonding can be described in terms of resonance between the structures.

Question 53

Definition Formal charge

Answer 53

the charge an atom would have if all covalent bonds in a molecule were broken and the electrons
assigned equally to the atoms involved.

Question 54

What does formal charge equal?

Answer 54

Formal charge = (no. of valence electrons in neutral atom)– (no. of lone pair electrons)– (1/2)(no. of electrons in bonds formed).

Question 55

VSEPR predictions of molecular shape centre on two fundamental requirements:

Answer 55

• Being able to assess the total number of electrons in the (sigma) bonding framework of a molecule
• Relating this number of electrons to a molecular shape.

Question 56

How to make a VSEPR prediction:

Answer 56

a) Write down Lewis structure for a molecule and identify the central atom

b) Count the number of σ bonding electron pairs and lone pairs around the central atom. To achieve the lowest energy these electron pairs position themselves so they are as far apart as possible:

From the arrangement of electron pairs, the location of the atoms and hence the shape of molecule can
be identified. These electron pairs fit in different electron domains (domain can be lone pair, single double etc).

Importantly, we need to keep the electron domains as far apart as possible and the lone pairs
repel more than bonding pairs.

Question 57

What are the energies of different electron pair repulsions?

(using inequalities)

Answer 57

Lone pair-lone pair > lone pair-bonding pair > bonding pair-bonding pair

This is because we need to keep the electron domains as far apart as possible and the lone pairs repel more than bonding pairs.

Question 58

What are the 6 basic geometries

Answer 58

linear
trigonal
tetrahedral
trigonal bipyramidal
octahedral
pentagonal bipyramid (rare)

In the trigonal bipyramidal geometry and the octahedral geometry there are equatorial (think of the middle / centre plan) and axial sites (ends).

In a trigonal bipyramidal arrangement the lone pairs occupy equatorial sites. In an octahedral
arrangement if there are two lone pairs these are positioned trans (or across the equatorial sites) to each other.

When lone pairs are
involved, there are
more geometries
that can occur

Question 59

What are electron domains?

Answer 59

an electron domain is a single
bond, a double bond a triple
bond or a lone pair

Question 60

Steric Number meaning
(VSEPR Geometries table)

Answer 60

the amount of electron domains

Question 61

Where to find VSEPR examples

Answer 61

week 3 slide 65 onwards

Question 62

What is the meaning of lone pairs of electrons?

Answer 62

In a Lewis structure representation of a molecule or ion, lone pairs are usually depicted as pairs of dots around the central atom.
They are valence electrons that don’t get involved in bonding.

Question 63

VSEPR exceptions

Answer 63

VSEPR is a useful tool that works well for most main group compounds (but not transition metal
compounds) although some exceptions are known.