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Chemistry Physical 1 > Chapter 3 (Bonding) > Flashcards

Chapter 3 (Bonding) Flashcards

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1
Q

Describe Ionic Bonding

A
  • Between metals and non-metals
  • Electrons are transferred from metal atoms to non-metal atoms.
  • Positive and negative ions are made
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2
Q

What are electrostatic forces?

A

The attraction between two oppositely charged ions.

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3
Q

What structure is formed by the electrostatic forces between ionic compounds?

A

Giant ionic lattice

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4
Q

What are the properties of ionically bonded compounds?

A
  • Always solid at room temperature
  • High melting point because of large structure and strong bonds.
  • Conduct electricity when molten or dissolved in water because the ions are then free to move.
  • Brittle and shatter easily as they form a lattice of alternating positive and negative ions which if moved together can create contact between charges.
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5
Q

Describe Covalent Bonding

A
  • A shared pair of electrons
  • Between two non-metals
  • The atoms share some of their outer shell electrons so that each atom has a stable noble gas arrangement.
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6
Q

How is a covalent bond drawn?

A

A line between two element symbols.

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7
Q

How do the shared electrons in covalent bonds hold atoms together?

A

Electrostatic attraction between nuclei and the shared electrons.

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8
Q

How many electrons are shared in a double covalent bond?

A

4

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9
Q

What are properties of substance with molecular structures?

A
  • Low melting temperatures as strong covalent bonds are only between the atoms within the molecule and weak attraction between the molecules so they don’t require much energy to move the molecules apart.
  • Do not conduct electricity because molecules are neutral overall so no charged particles to carry current.
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10
Q

What is co-ordinate/dative covalent bonding?

A

Where one atom provides both electrons in a covalent bond.

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11
Q

How does co-ordinate/dative covalent work?

A
  • The atom that accepts the electron pair is electron-deficient (doesn’t have a full outer shell)
  • The atom that is donating the electrons has a pair of electrons which is not being used in a bond, called a lone pair.
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12
Q

Describe Metallic Bonding

A

As metal atoms cannot transfer atoms to another metal the main outer shell of atoms merge so that they are no longer associated with a particular atom so a lattice of positive ions and a ‘sea’ of delocalised electrons is formed.

The positive ions tend to repel each other but this is balanced by the negatively charged delocalised electrons.

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13
Q

Why can metals conduct electricity?

A

The delocalsied electrons can move through the structure and an electron from a negative terminal of the supply joins the electron at one end of the wire while at the same time a different electron leaves the wire at the positive terminal.

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14
Q

Why can metals conduct heat?

A

The sea of delocalised electrons is partially responsible and the closely packed ions can also spread energy by vibrations.

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15
Q

Why are metals strong?

A

The delocalised electrons extend throughout the solid so there are no individual bonds to break.

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16
Q

What effects the strength of a metal?

A
  • The charge on the ion - The greater the charge on the ion, the greater the number of delocalised electrons and the stronger the attraction between the positive ions and the electrons.
  • The size of the ion - The smaller the ion the closer the electrons are to the positive nucleus so the stronger the bond.
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17
Q

Why are metals malleable and ductile?

A

The positive ions are arranged in layers that can slide over each other.

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18
Q

Why do metals have high melting points?

A

They have giant structures, there is a strong attraction between positive metal ions and delocalised sea of electrons.

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19
Q

Define electronegativity

A

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

20
Q

What factors effect electronegativity?

A
  • Nuclear charge
  • The distance between the nucleus and the outer shell electrons.
  • The shielding of the nuclear charge by electrons in inner shells.
21
Q

How does electronegativity change going down a group?

A

Electronegativity decreases as there is more shielding as there are more inner shells.

22
Q

How does electronegativity change going across group?

A

Increases as nuclear charge increases and the number of inner shells remain the same. Atom contracts and gets smaller due to nuclear charge.

23
Q

What is the polarity of covalent bonds?

A

The unequal sharing of electrons between atoms that are bonded together covalently.

24
Q

Describe the polarity in a covalent bond with two of the same elements?

A

The atoms are non-polar as they have the same electronegativity. The electrons in the bond are shared equally between the atoms.

25
Q

Describe the polarity in a covalent bond with two different elements?

A

The electrons will be more attracted to the more electronegative molecule so will not be distributed equally around the molecule.

26
Q

What are the 3 types of intermolecular forces in order of weakest to strongest?

A

van der Waals, dipole-dipole, hydrogen

27
Q

Where do dipole-dipole forces occur?

A

Between 2 molecules that have a permanent dipole.

28
Q

What is a dipole?

A

The effect of the polarity of all bonds in the molecules.

29
Q

What are van der Waals forces?

A

The very weak electrostatic forces between all atoms and molecules (as they contain both positive and negative charges, even though they are neutral overall)

30
Q

How do van der Waals forces work?

A

The negative electrons in an atom can be in any position at one time so the distribution of charges changes every instant. This creates a temporary dipole that which effects electron distribution in near-by atoms. As electron distribution of original atom changes, it creates an induced dipole force which is called a van der Waals force.

31
Q

What factor increases the effect of van der Waals forces?

A

Number of electrons present (larger atomic or molecular mass)

32
Q

Which elements are able to create hydrogen bonds?

A

Fluorine, Oxygen, Nitrogen

33
Q

Describe how hydrogen bonds form using water as an example

A

1) Oxygen has a lone pair of electrons.
2) in water the hydrogen atoms are highly electron-deficient as oxygen is very electronegative and attracts the electrons in the bond towards it.
3) The hydrogen atoms in the water are positively charged and very small. They expose protons which have a very strong electric field because of their small size.

34
Q

What two factors must be included to create hydrogen bonds?

A
  • A hydrogen atom that is bonded to a very electronegative atom. This will produce a strong partial charge on the hydrogen atom.
  • A very electronegative atom with a lone pair of electrons. These will be attracted to the partially charged hydrogen atoms in another molecule and for the bond.
35
Q

How much stronger are covalent bonds than hydrogen?

A

10 times

36
Q

How do hydrogen bonds change as water is turned from liquid to ice?

A

In a liquid, the hydrogen bonds are broken and reformed as the molecules are able to move around. When the water freezes then the molecules are no longer free to move as the hydrogen bonds hold the molecules in fixed positions.

37
Q

What is electron pair repulsion theory?

A

The idea that each pair of electrons around an atom will repel all other electron pairs and the pairs of electrons will, therefore, take up positions as far apart as possible to minimise repulsion.

38
Q

What holds ions together in an ionic lattice?

A

Electrostatic attraction (attraction between oppositely charged particles)

39
Q

Why do ionic compounds have high melting points?

A

They have strong electrostatic attractions that extend throughout the structure and requires and a lot of energy to break as well as a large surface area.

40
Q

Describe a metallic crystal

A

A lattice of positive ions which are embedded in a delocalised sea of electrons. The attraction of positive ions and negative electrons extends throughout the structure which results in the high melting temperature.

41
Q

Describe molecular crystals

A

Consist of molecules held in regular array by intermolecular forces (van der Waals). Have covalent bonds within the molecules holding the atoms together but not between the molecules. Intermolecular forces are weak so so molecular crystals have low melting points.

42
Q

What are the properties of molecular crystals?

A
  • Soft and easily breakable
  • Low melting temperatures and sublimes readily form gaseous molecules
  • Do not conduct electricity because there are no charged particles to carry a charge
43
Q

Describe the structure of diamond

A

Consists of pure carbon with covalent bonding between each carbon atom bonds spread throughout the structure to make it a macromolecular structure.

In diamond each carbon bonds with 4 other carbons so four electron pairs repel each other.

44
Q

What are the properties of diamond?

A
  • Very hard material
  • Very high melting temperature (over 3700K)
  • Does not conduct electricity as there are no free charged particles to carry a charge.
45
Q

Describe the structure of graphite

A

Pure carbon with both strong covalent bonds between atoms and weaker van der Waals forces between layers. Each carbon has 3 covalent bonds so there a ‘p’ orbital with a spare electron for each carbon which acts as a delocalised electron so can carry a charge.

46
Q

What are the properties of graphite?

A
  • Very soft with layers that can slide over each other.
  • Very high melting point due to strong covalent bonds.
  • Conducts electricity

Chemistry Physical 1 (8 decks)

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