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Chemistry > Atomic Structure and Bonding > Flashcards

Atomic Structure and Bonding Flashcards

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1
Q

Periods

A

Horizontal rows on the periodic tabl. All elements in the same period have the same number of electron shells.

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2
Q

Groups

A

Vertical columns on the periodic table. All elements in the same group have similar chemical properties and the same numbe rof valence electrons.

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3
Q

Mass of a Proton, Neutron and Electron.

A

1, 1, 0.0005 (negligable)

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4
Q

Charge of a Proton, Neutron and Electron

A

1+, 0, 1-

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5
Q

What is the top number of an element on the periodic table?
Eg: 3
Li
lithium
7

A

Atomic number.
(also number of protons/electrons in the atom)

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6
Q

What is the bottom number of an element on the periodic table?
Eg: 3
Li
lithium
7

A

Atomoic mass number.
(also number of protons + neutrons)

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7
Q

Octet Rule

A

The tendency of atoms to prefer to have 8 electrons in the valence shell. Rule: [2,8,8,2] or [2,8,18,18,8,2]

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8
Q

Halogen

A

All elements in group 7.

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9
Q

Alkali metals

A

All elements in group 1.

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10
Q

Noble Gases

A

All elements in group 8. Their valence shells are full so they cannot become bonded.

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11
Q

Ionic Bond

A

The electrostatic force of attraction between oppositely charged ions (cations and anions / metal and non-metal).

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12
Q

Lattice

A

A regular repeated 3-D arrangement

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13
Q

Giant

A

Undefined number of particles (eg: ions) in the lattice. Simplest whole number ratio.

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14
Q

Polyatomic ion

A

A chemical ion composed of 2 or more atoms covalently bonded. (eg: SO4)

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15
Q

Common polyatomic ions (names AND formulas)

A

Carbonate: CO3 (2-)
Sulfate: SO4 (2-)
Hydroxide: OH (-)
Nitrate: NO3 (-)
Phosphate: PO4 (3-)
Hydrogen Carbonate: HCO3 (-)
Ammonium: NH4 (+)
(only positively charged polyatomic ion)

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16
Q

Properties of ionic compounds

A
  • High melting and boiling point
  • Hard
  • Brittle
  • Often soluble in water
  • Doesn’t conduct electricity as a solid
  • Can conduct electricity when molten or dissolved
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17
Q

How do ions have high melting and boiling points?

A

The strong electrostatic forces of attraction between the ions require a lot of heat energy to break.

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18
Q

How are ions ‘mostly soluble in water’?

A

If an ionic compound is soluble, water molecules are able to break the electrostatic force of attraction between the ions.

19
Q

Describe conductivity in solid and liquid ions.

A

Ions in a solid ionic compound are not free to move and therefore cannot conduct electricity.
Whereas in a liquid/solution, ions move past eachother and charge can flow, inducing a current.

20
Q

How are ions ‘brittle’?

A

When a force is applied, the ions in the lattice shift so that like charges repel and the lattice breaks.

21
Q

How are ions ‘hard’?

A

Ionic compounds have strong electrostatic forces of attractions between oppositely charged ions.

22
Q

Flame tests

A

Flame tests are used to detect the presence of a metal ion.
Lithium (Li+) = red flame
Sodium (Na+) = intense yellow/orange flame
Potassium (K+) = purple/lilac/violet flame
Calcium (Ca 2+) = brick red flame

23
Q

Ion Test - Ammonium (NH4 +)

A
  1. add dilute sodium hydroxide solution
  2. gently heat
  3. if NH4 + is present: choking smell produced, fumes turn damp red litmus paper or damp universal indicator from red to blue
24
Q

Ion Test - Carbonate (CO3 2-)

A
  1. add dilute hydrochloric acid
  2. bubble gas through limewater
  3. if CO3 2- is present, limewater turns cloudy
25
Q

Ion Test - Sulfate (SO4 2-)

A
  1. add dilute hydrochloric acid
  2. add a few drops of barium chloride solution
  3. if SO4 2- is present, a white precipitate forms
26
Q

Ion Test - Hallide ions
Chloride ions (Cl-)
Bromide ions (Br-)
Iodide ions (I-)

A
  1. add dilute nitric acid
  2. then add a few drops of silver nitrate solution
  3. if Cl- is present, a white precipitate forms
    if Br- is present, a cream precipitate forms
    if I- is present, a yellow precipitate forms
27
Q

General properties of metals

A
  • High melting and boiling point
  • Good conductors of heat
  • Hard
  • High density
  • Malleable
  • Ductile
  • Sonorous
  • Lustrous

(exceptions: Sodium is less hard and can be cut with a knife, Mercury has a low melting and boiling point and is liquid at room temp)

28
Q

Structure/bonding of metals

A

Metal cations are attracted to the sea of delocalised electrons. This electrostatic attraction is metallic bonding. The atoms in a pure metal are in tightly packed layers, which forms a giant metallic lattice.

29
Q

How are metals able to conduct electricity?

A

Delocalised electrons can move through the structure and carry charge.

30
Q

How are metals malleable?

A

The arrangement of ions can change but the attraction between cations and delocalised electrons isn’t broken.

31
Q

How are metals able to conduct heat?

A

Delocalised valence electrons can move faster and pass on energy quickly. The lattice of atoms transmits energy efficiently as the cations vibrate and bump into other cations.

32
Q

How do metals have high melting and boiling points?

A

The electrostatic force of attractions between cations and delocalised valence electrons are strong and it requires a lot of heat energy to break them.

33
Q

Alloys

A

A mixture composed of 2 or more elements, at least one of which is a metal. Alloys are generally harder, less malleable stronger and poorer conductors than pure metals because the arrangement of ions is disturbed.
Eg: Stainless steel, bronze and brass.

34
Q

Covalent bonding

A

A chemical bond that involves the sharing of electron pairs between atoms. Only involves valence electrons and is usually between non-metals. These bonds are very strong and form simple molecular structures.

35
Q

Properties of simple molecular structures

A
  • low melting and boiling points
  • can’t conduct electricity in any form
36
Q

How do simple molecular structures have low melting and boiling points?

A

There are intERmolecular forces between simple molecules which are weaker than the strong covalent bonds. When simple molecular substances melt or boil, its the weak intermolecular forces that are overcome - covalent bonds are not broken.

37
Q

How are simple molecular strustures not able to conduct electricity?

A

Simple molecules have no overall charge or charged particles that can seperate so they cannot conduct electricity.

38
Q

Polarity of simple molecular structures

A

Negative electrons are held in their shells by the attraction of the positive protons in the nucleus (the nucleus has a pull on electrons). Electrons shared in covalent bonds are pulled closer to the atom with a stronger nuclear pull. This creates a slight negative and a slight positive side in the bond. (Elements NOF have the strongest pull)

39
Q

Non-polar vs Polar

A

Non-polar = even distipution of charge, symmetrical. (eg: Methane / CH4)
Polar = uneven distribution of charge, unsymmetrical (Hydrogen Fluoride / HF)

40
Q

Giant covalent structures + properties

A

Contain many (undefined number) of atoms, each joined to adjacent atoms by covalent bonds. These atoms are usually arranged into a giant regular lattice, extremely strong structures.
Properties= very high melting points and usually can’t conduct electricity (graphite is the exception)

41
Q

Silicon dioxide / Silica

A

Main component found in sand. All atoms (silicon and oxygen) are covalently bonded in a regular arrangement forming a giant covalent structure.

42
Q

Allotropes of carbon

A

Diamond and Graphite

43
Q

Diamond

A

Is a giant covalent structure and allotrope of carbon. Is extremely hard and has a high melting point as all carbon atoms are covalently bonded to 4 other carbon atoms - leaving no free electrons and no ions. Is insoluble in water and cannot conduct electricity.

44
Q

Graphite

A

Is a giant covalent structure and allotrope of carbon. Is insoluble in water, has a high melting point and is a good conductor of electricity. Each carbon atom is only covalently bonded to 3 other carbon atoms, leaving each atom with a spare elctron, forming a sea of delocalised electrons which can carry charge.
Graphite is made of many layers joined by intermolecular forces, making it easy to break these layers and move around eachother, making graphite slippery.

Chemistry (2 decks)

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