Bonding

Question 1

Describe the shape of two pairs of electrons around a central atom (draw an example of this)

Answer 1

linear= 180 degrees apart

Question 2

Describe the shape of three pairs of electrons around a central atom (draw an example of this)

Answer 2

Trigonal planar = 120 degrees

Question 3

Describe the shape of four pairs of electrons around a central atom (draw an example of this)

Answer 3

• tetrahedral = 109.5 degrees
• 3 dimensional not planar = arrangement of the angles can be more than 360 degrees

Question 4

Describe the shape of five pairs of electrons around a central atom (draw an example of this)

Answer 4

trigonal bipyramid

Question 5

Describe the shape of six pairs of electrons around a central atom (draw an example of this)

Answer 5

octahedral = 90 degrees

Question 6

Draw a molecules with 2 bonding pairs and one lone pair. Name the structure and the angles in between.

Answer 6

Question 7

Draw a molecule with 3 bonding pairs and one lone pair. Name the structure and the angles in between.

Answer 7

107 degrees, trigonal pyramid

Question 8

Draw a molecule with 2 bonding pairs and 2 lone pair. Name the structure and the angles in between.

Answer 8

104.5 = bent planar

Question 9

Draw a molecule with 4 bonding pairs and 1 lone pair. Name the structure and the angles in between.

Answer 9

89 and 119 = seesaw

Question 10

Draw a molecule with 3 bonding pairs and 2 lone pair. Name the structure and the angles in between.

Answer 10

89 degrees = t shape

Question 11

Draw a molecule with 2 bonding pairs and 3 lone pair. Name the structure and the angles in between.

Answer 11

Question 12

Draw a molecule with 5 bonding pairs and 1 lone pair. Name the structure and the angles in between.

Answer 12

89 degrees

Question 13

Draw a molecule with 4 bonding pairs and 2 lone pair. Name the structure and the angles in between.

Answer 13

Question 14

Draw a molecule with 3 bonding pairs and 3 lone pair. Name the structure and the angles in between.

Answer 14

89 degrees

Question 15

Draw a molecule with 2 bonding pairs and 4 lone pair. Name the structure and the angles in between.

Answer 15

Question 16

What is ionic bonding

Answer 16

the result of electrostatic attraction between oppositely charged ions

Question 17

Name and explain the chemistry behind the properties of ionic compounds/crystals

Answer 17

• Solids at room temperature and strong = They have giant structures, lots of energy needed to break up the lattice of ions. There are strong electrostatic attractions between ions.
• conduct electricity when molten or dissolved in water = ions (charged particles) are free to move throughout the structure when molten, so charge is also able to move throughout the structure.
• Brittle and shatter easily= they form a lattice of alternating positive and negative ions , so when they are hit, the ions may move and have contact between ions with the same charge, and they repel
• soluble = water is polar, so it attracts oppositely charged ions.

Question 18

What is covalent bonding?

Answer 18

• covalent bond forms between a pair of non-metal atoms
• the atoms share some of their outer electrons so they each atom has a stable noble gas arrangement
• shared pair of electrons

Question 19

How does sharing electrons hold atoms together?

Answer 19

held together by the electrostatic attracting between the nuclei and the shared electrons

Question 20

What are the properties of substances with molecular structures?

Answer 20

• low melting temperatures = strong covalent bonds are only between atoms within the molecules. there is only a weak attraction between the molecules do not need much energy to move apart from each other.
• poor conductors of electricity = the molecules are neutral overall, so they are no charged particles to carry the current. this happens even if they dissolve in water

Question 21

What is co-ordinate bonding?

Answer 21

when one atom provides both the electrons

• the atom that accepts the electron pair is an atom that does not have a filled outer main level of electrons ( electron deficient)
• the atom that is donating the electrons have a pair of electrons that is not being usied in a bond (lone pair)

Question 22

Why are metals good conductors of electricity and heat

Answer 22

• delocalized electrons can move throughout the structure = allows charge to flow throughout the structure
• sea of electrons means that heat energy is spread by increasingly vigorous vibrations of the closely packed ions

Question 23

What factors affect the strength of metals

Answer 23

• the charge on the ion = the greater the charge on the ion, the greater the number of delocalized electrons and the stronger the electrostatic attraction been the positive ions and the electrons
• the size of the ion - the smaller the ion, the closer the electrons are to the positive nucleus and the stronger the bond
• delocalized electrons = these extend throughout the solid so there are o individual bonds to break.

Question 24

Why do metals have high melting points?

Answer 24

because they have giant structures. there is a strong attraction between metal ions and the delocalized sea of electrons, making atoms difficult to separate.

Question 25

What is electronegativity?

Answer 25

the power of an atom to attract the electron density in a covalent bond towards itself

Question 26

What is electron density?

Answer 26

the way the negative charge is distributed in a molecule

Question 27

Describe the Pauling scale

Answer 27

- used as a measure of electronegativity - runs from 0 to 4 - the greater the number = more electronegativity

Question 28

Why are noble gases 0 on the Pauling scale?

Answer 28

because, in general, they do not form covalent bonds with other atoms

Question 29

What does electronegativity depend on?

Answer 29

- the nuclear charge (more protons) = greater attraction to shred pair of electrons - the distance between the nucleus and the outer shell electrons = closer means stronger attraction - the shielding of the nuclear charge by electrons in the inner shells= less shielding, greater attraction to shared pair of electrons - The smaller the atom, the closer the nucleus is to the share outer main level electrons and the greater its electronegativity

Question 30

Describe the trend in electronegativity as you go up a group in the Periodic Table

Answer 30

electronegativity increases, due to the atoms getting smaller and there is less shielding by electrons in inner shells

Question 31

Describe the trend in electronegativity as you go across a period in the Periodic Table

Answer 31

Increases = the nuclear charge increases, the shielding is constant and the atom becomes smaller

Question 32

What is polarity?

Answer 32

the unequal sharing of the electrons between atoms that are bonded together covalently

Question 33

Describe the electronegativity in a pure covalent bond

Answer 33

In a pure covalent bond, the bond must be shared equally between atoms = both atoms have exactly the same electronegativity and the bond is completely non-polar

Question 34

Describe covalent bonds between two atoms that are of a different electronegativity

Answer 34

the electrons in the bond will not be shared equally between the atoms = polar molecule. = polar bond due to s+ and s- poles ## Footnote **unsymmetrical electron distribution**

Question 35

Why are some molecules with polar ends not a polar molecule

Answer 35

because sometimes the electrons density poles pull in opposite directions

Question 36

What are the three intermolecular forces?

Answer 36

Van der Waals forces Dipole-dipole forces Hydrogen bonding

Question 37

What is a dipole?

Answer 37

a difference in charge between the two atoms caused by a shift in electron density in the bond

Question 38

What a dipole-dipole forces?

Answer 38

act between molecules that have permanent dipole= weak electrostatic forces of attraction between s+ and s- charges on neighbouring molecules

Question 39

What are van der Waals forces?

Answer 39

weak electrostatic attractions that occur between positive and negative charges: happens between all atoms because they are made up of positive and negative charges. they are in addition to any other intermolecular forces

Question 40

How do van der Waals forces work?

Answer 40

**Electrons (in charge clouds) are always moving really quickly**, hence their distribution is sometimes not even = **temporary dipole.** This dipole can **cause another temporary dipole i**n the opposite direction on a neighboring atom. **these two dipoles are then attracted to each other.** So on and so on (domino effect) Dipoles are constantly created and destroyed, however, **the overall effect is for the atoms to be attracted to each other.** The more electrons there are, the larger the temporary dipole

Question 41

What increases the strength of van der Waals?

Answer 41

the number of electrons (higher = stronger) larger atomic masses = strong van der Waals = the number of electrons increases, and so also does the radius of the atom. The more electrons you have, and the more distance over which they can move, the bigger the possible temporary dipoles and therefore the bigger the dispersion forces.

Question 42

What is hydrogen bonding?

Answer 42

when a hyrogern atom (s+) interacts with a more electronegative atom with a s- charge (oxygen, fluorine, nitrogen are the only elements electronegative enough), attracted to lone pairs (but doesn't form dative bonds)

Question 43

Why is hydrogen bonding so strong?

Answer 43

Hydrogen atoms become very electron deficient when bonding to fluorine, nitrogen and oxygen. This is because these elements are very electronegative and attracts the shared electrons in the bond towards it. Hydrogen (s+) atoms are very small. These exposed protons have a very strong electric field because of heir small size. Lone pairs also become strongly attracted to the electron deficient hydrogen

Question 44

What is the electron pair repulsion theory?

Answer 44

- each pair of electrons around an atom ill repel all other electron pairs - the pairs of electrons will therefore take up a position as far apart as possible to minimize repulsion

Question 45

Describe what happens when you heat a solid?

Answer 45

energy is supplied to the particles = makes them vibrate more about a fixed position = slightly increases the average distance between the particles and so the solid expands

Question 46

What happens during fusion?

Answer 46

A solid is supplied with so much energy that the forces that act between the solid particles are weakened = liquid. the energy required for this to happen i called enthalpy change of melting

Question 47

Why doesn't the temperature change whiles the solid is melting?

Answer 47

because the heat energy provided is absorbed as the forces between particles are weakened.

Question 48

What happens during vaporization?

Answer 48

lots of energy is supplied to the liquid to break all of the intermolecular forces between the particles (as gas move freely and independently). The energy required to do this is called the enthalpy change of vaporisation

Question 49

What happens when you heat a gas?

Answer 49

the particles gain kinetic energy and move faster. they get much further apart (expand greatly). this increases the pressure when in a container.

Question 50

Describe the bonding and properties of molecular crystals (iodine)

Answer 50

- covalent bonds within the molecules and intermolecular forces between them - low melting/boiling point and brittle: weak van der Waals forces - poor electrical conductivity= no ions to conduct and electrons localised - poor solubility = iodine is non-polar. Water is polar, so water molecules remain hydrogen-bonded with each other, rather than forming van der Waals with iodine.

Question 51

What is the structure and what are the properties of diamond?

Answer 51

4 single covalent bonds between carbon atoms very high melting point = thousands of strong covalent bonds between carbon atoms= lots of energy needed to overcome no electrical conductivity = no delocalized electrons, all electrons used in covalent bonds very hard and strong = structure very rigid, arrangement of atoms held by covalent bonds insoluble= no charged particles, so water is not attracted to it. water is not strong enough to break covalent bonds.

Question 52

What is the structure and what are the properties of graphite?

Answer 52

- each carbon atom forms three single covalent bonds to other carbon atoms - very high melting and boiling points = thousands of strong covalent bonds between carbon atoms in the layers = lots of energy needed - very good electrical conductivity= delocalized electrons can move throughout the structure, so charge can flow through the structure - very strong = due to its unbroken pattern and the strong covalent bonds between carbon atoms - flexible= strong bonds between carbon atoms are flexible, can be twisted without breaking

Question 53

What is the property of ice?

Answer 53

low melting and boiling point: hydrogen bonds between molecules are weak no electrical conductivity: no complete ions to allow charge through structure less dense: in ice, water molecules have 4 hydrogen bonds and are in a rigid structure. In water, hydrogen bonds are being constantly made and broken, allowing the water molecules to be closer to each other= more water molecules per cm^3= more dense

Question 54

**Why is graphite soft?**

Answer 54

weak forces between planes

Question 55

**Suggest why lone pairs are sometimes opposite from each other**

Answer 55

lone pairs repel more, so they are the furthest away from each other

Question 56

What is the formula for the compound ion sulfate?

Answer 56

SO₄^2-

Question 57

What is the formula for the compound ion hydroxide ?

Answer 57

OH^-

Question 58

What is the formula for the compound ion nitrate?

Answer 58

NO₃-

Question 59

What is the formula for the compound ion carbonate?

Answer 59

CO₃^-2

Question 60

What is the formula for the compound ion ammonium?

Answer 60

NH₄⁺