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Year 10 Chemistry > Periodic Chemistry > Flashcards

Periodic Chemistry Flashcards

1
Q

What is periodic chemistry?

A

Periodic chemistry involves the study of how physical and chemical properties of elements vary with their different electron arrangements.

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2
Q

Describe the Alkali Metals in Group 1.

A

The Alkali Metals in Group 1 are Lithium, Sodium, Potassium, Rubidium, Cesium, and Francium. They are univalent, highly reactive metals that easily lose electrons to become positively charged ions.

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3
Q

What are some uses of helium and argon?

A

Helium is used in filling meteorological balloons and airships, while argon is used to provide an inert atmosphere for welding.

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4
Q

What is the main characteristic of noble gases?

A

Noble gases are generally chemically inactive due to their stable completed outer shell electronic configuration of 2 or 8.

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5
Q

What is the common oxidation number of halogens?

A

The common oxidation number of halogens is -1.

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6
Q

Describe the properties of halogens.

A

Halogens are fluorine, chlorine, bromine, iodine, and astatine. They exist as diatomic molecules, are highly reactive, and can easily gain one electron to form anions with an oxidation number of -1.

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7
Q

How do group 2 elements compare to group 1 metals in terms of reactivity, melting points, and hardness?

A

Group 2 elements are less reactive than group 1 metals, have higher melting points, and are harder.

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8
Q

What is the oxidation state of group 2 elements?

A

The oxidation state of group 2 elements is +2.

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9
Q

Describe the group 2 elements.

A

The group 2 elements include beryllium, magnesium, calcium, strontium, barium, and radium. They have a fixed oxidation state of +2 and ionize by giving off 2 electrons

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10
Q

What are the properties of Alkaline Earth Metals?

A

Alkaline Earth Metals are malleable, ductile, and good conductors of heat and electricity. They have higher melting points and densities compared to Alkali Metals.

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11
Q

What are the properties of Alkali Metals?

A

Alkali Metals are strongly electropositive, easily tarnish in air, are good conductors of heat and electricity, and have low melting points and densities.

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12
Q

Define a period in the periodic table.

A

A period in the periodic table is a set of elements with the same number of electron shells, where atomic numbers increase by one unit from one atom to the next.

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13
Q

What is a group in the periodic table?

A

A group in the periodic table consists of elements with the same number of valence electrons in their outermost shells.

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14
Q

Describe the periodic table.

A

The periodic table is a table showing all known elements arranged systematically in order of increasing atomic number, with elements in the same group having the same number of valence electrons.

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15
Q

What does the periodic law state?

A

The periodic law states that the properties of elements are a periodic function of their atomic numbers.

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16
Q

What is the formula to determine the maximum number of electrons in a shell?

A

The formula is 2n^2, where n is the shell number.

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17
Q

What is the maximum number of electrons in an outermost shell?

A

The maximum number of electrons in an outermost shell is 8.

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18
Q

Explain the concept of core electrons.

A

Core electrons, or inner electrons, are electrons between the nucleus and the valence electrons. They do not participate in chemical reactions but influence the behavior of valence electrons.

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19
Q

What are valence electrons?

A

Valence electrons are the electrons in the outermost shell of an atom, responsible for the chemical properties of elements.

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20
Q

Describe the arrangement of electrons in shells.

A

Electrons in an atom can orbit in a maximum of seven shells or energy levels, represented by letters K, L, M, N, O, P, and Q. Each shell has different energies.

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21
Q

Define atomic number.

A

Atomic number is the number of protons in the nucleus and is equal to the number of electrons in a neutral atom.

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22
Q

Describe transition metals

A

Transition metals are elements that have partially filled d orbitals

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23
Q

List the atomic numbers for the first row transition metals.

A

The atomic numbers for the first row transition metals range from 21 to 30.

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24
Q

What are some general properties of transition metals?

A

Transition metals are hard with high melting points, conduct electricity, and are malleable and ductile

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25
Q

What are some characteristics of transition metals?

A

Transition metals exhibit variable oxidation states, form colored hydrated salts or ions, are often used as catalysts, form complexes, and show paramagnetic properties.

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26
Q

Define metalloids.

A

Metalloids are elements that have a mixture of metallic and non-metallic properties

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27
Q

List some physical properties of metalloids.

A

Metalloids are brittle and lustrous, they are not good conductors but can become good conductors when slightly impure, and they are good insulators.

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28
Q

What are some examples of metalloids?

A

Examples of metalloids are Germanium (Ge), Boron (B), and Silicon (Si).

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29
Q

Describe metallic character.

A

Metallic character refers to the ability of metals to lose electrons from their outermost shells to become cations.

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30
Q

What are some physical properties of metals?

A

Metals are lustrous, malleable and ductile, good conductors of heat and electricity, and hard with high melting points.

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31
Q

What are non-metals?

A

Non-metals are elements with four or more electrons in the outermost shells of their atoms.

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32
Q

Describe the physical properties of non-metals.

A

Non-metals are insulators and do not conduct heat and electricity, except for graphite. They are also brittle.

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33
Q

List some examples of non-metals.

A

Examples of non-metals include phosphorus, sulfur, iodine, oxygen, and nitrogen.

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34
Q

What are periodic properties?

A

Periodic properties are the periodic variations in physical or chemical properties of elements with increasing atomic number.

35
Q

Define atomic radius.

A

Atomic radius is half the distance between the nuclei of any two similar close atoms in a substance.

36
Q

What are the three types of atomic radii?

A

The three types of atomic radii are covalent radius, metallic radius, and van der Waals radius.

37
Q

Explain effective nuclear charge.

A

Effective nuclear charge is the net positive charge that the outermost electrons feel through attraction from protons in the nucleus.

38
Q

What is the shielding effect?

A

The shielding effect involves the repulsion between inner filled orbitals/core electrons and the electrons being removed from the outer orbital, reducing the total nuclear attraction.

39
Q

How does atomic radius vary with effective nuclear charge?

A

Atomic radius decreases with increasing effective nuclear charge.

40
Q

Describe the screening effect in atomic size.

A

The screening effect is the reduced attraction by the protons for the outermost electrons due to the intervening electrons.

41
Q

Explain the trend in atomic size across a period.

A

Across a period, atomic size decreases. This is because as atomic number/nuclear charge increases, the effective nuclear charge also increases, causing the attraction for the outermost electrons by the nucleus to increase and the electrons to be drawn closer to the nucleus.

42
Q

Describe the trend in atomic size down a group.

A

Down a group, atomic size increases. This is because as the atomic number increases, the principal quantum number of the valence electron also increases, resulting in the addition of a new shell. This increases the screening effect, decreases the effective nuclear charge, and causes the outermost electrons to be less attracted to the nucleus.

43
Q

What is the diagonal relationship in the periodic table?

A

The diagonal relationship refers to similarities that exist between pairs of elements in different groups and periods of the periodic table. These pairs of elements have similar atomic and ionic radii, as well as similar chemical properties, due to similar charge densities.

44
Q

Define ionic radius.

A

Ionic radius describes the size of an ion in a crystal. It is determined by the addition or subtraction of electron(s) from an atom.

45
Q

Explain the trend in cation size across a period.

A

Across a period, the size of cations decreases. This is because when an electron is removed to form a cation, it results in a decrease in electron repulsions.

46
Q

Describe the trend in cation size down a group.

A

Down a group, the size of cations increases. This is because the addition of a new shell as the atomic number increases increases the screening effect, decreases the effective nuclear charge, and causes the outermost electrons to be less attracted to the nucleus.

47
Q

How do cations compare in size to their corresponding atoms?

A

Cations are smaller than their corresponding atoms. For example, Na+ is smaller than Na.

48
Q

Describe the effect of effective nuclear charge on the size of an ion.

A

The effective nuclear charge increases, resulting in a strong attraction of the outermost electrons towards the center and decreasing the size of the ion.

49
Q

Explain why anions are generally larger in size than their corresponding atoms.

A

Anions are formed by the gain of electron(s), which increases the repulsion and decreases the effective attraction of the nucleus for the outermost electron, causing the electrons to spread out and increase the size of the anion.

50
Q

Define isoelectronic series.

A

Isoelectronic series is a set of atoms and ions with the same electronic configuration.

51
Q

What factors determine the size of species in an isoelectronic series?

A

The size of species in an isoelectronic series depends on their nuclear charge (number of protons) and the size of their positive nuclear charge if they have the same type of charge.

52
Q

Arrange the following species in order of increasing size: F-, Na+, Mg2+, Al3+. Provide a reason for your order.

A

The order of increasing size is F- < Na+ < Mg2+ < Al3+. This is because the number of protons (nuclear charge) increases in that order, resulting in an increasing inward pull of nuclear charge on the valence electrons and decreasing size. Anions are generally bigger than cations due to repulsion causing electrons to spread out.

53
Q

Describe ionization energy.

A

Ionization energy is the minimum energy required to remove an electron from the outermost orbital of a gaseous isolated atom or ion.

54
Q

What is the first ionization energy?

A

The first ionization energy is the minimum amount of energy required to remove one electron from the outermost orbital of an atom to form one mole of gaseous isolated cation with a single charge.

55
Q

Define second ionization energy.

A

The second ionization energy is the minimum amount of energy required to remove an electron from one mole of a singly charged cation in the gaseous state to form one mole of a cation with a double charge.

56
Q

How is total ionization energy calculated?

A

the total ionization energy is the sum of the ionization energies.

57
Q

What factors affect ionization energy?

A

The factors affecting ionization energy are effective nuclear charge, atomic size, and penetration effect.

58
Q

Describe the relationship between the size of the positive charge on an ion and its ionization energy.

A

The higher the positive charge on the ion, the higher the ionization energy due to increased effective nuclear charge and the need to break into a new orbital.

59
Q

Explain how the stability of the electronic configuration affects ionization energy.

A

More stable configurations make it more difficult to remove electrons, resulting in higher ionization energy.

60
Q

Define atomization energy

A

Atomization energy is the energy absorbed when one mole of free gaseous atoms is formed from an element at standard conditions.

61
Q

Describe the type of bond that can be predicted when the electronegativity difference is less than or equal to 0.5

A

When the electronegativity difference is less than or equal to 0.5, a purely covalent bond is predicted.

62
Q

What type of compound is formed when the electronegativity difference is between 0.5 and 1.7?

A

A polar covalent compound is formed when the electronegativity difference is between 0.5 and 1.7.

63
Q

Arrange the compounds PH3, GeCl4, and CsF in decreasing order of ionic character.

A

CsF, GeCl4, PH3 (CsF has the highest ionic character, followed by GeCl4, and PH3 has the lowest ionic character).

64
Q

What is the electronegativity difference between P and H in PH3?

A

The electronegativity difference between P and H in PH3 is 0.

65
Q

Do electronegativity values increase or decrease across the periodic table?

A

Electronegativity values generally increase across the periodic table.

66
Q

Describe how electronegativity difference is used to determine the type of compound formed.

A

The change in electronegativity values can be used to determine whether a compound is highly ionic, polar covalent, or purely covalent.

67
Q

What does the electronegativity difference between atoms in a covalent compound indicate?

A

The electronegativity difference indicates the ionic character of the covalent bond or the overall ionic nature of the compound.

68
Q

Define electronegativity.

A

Electronegativity is the ability of an atom to attract electrons towards itself in a chemical bond.

69
Q

How does atomization energy change with increasing size of atoms in molecules of the same type of bonds?

A

Atomization energy decreases with increasing size of the atoms

70
Q

Describe the relationship between bond strength and atomization energy.

A

When a bond is very strong, its atomization energy is very high.

71
Q

Describe the process of sublimation.

A

Sublimation occurs when a solid absorbs heat and changes directly into the gaseous state without first changing into a liquid.

72
Q

What is the relationship between atomization energy and sublimation energy?

A

Atomization energy is the same as sublimation energy for solid substances.

73
Q

Describe the trend of ionization energy across a period.

A

Ionization energy increases from left to right across a period due to increasing effective nuclear charge and decreasing size.

74
Q

does electronegativity change across a period?

A

Electronegativity increases across a period from left to right as the atomic size decreases and the effective nuclear charge increases.

75
Q

Why does fluorine have unexpectedly lower electron affinity compared to other Group 7 elements?

A

Fluorine has unexpectedly lower electron affinity due to its extremely small size, which creates electron-electron repulsion between its existing electrons and any incoming electron.

76
Q

How does electron affinity change down a group?

A

Electron affinity decreases down a group because the atomic radius increases, resulting in less energy being released when an electron is added.

77
Q

What factors influence electron affinity?

A

Factors influencing electron affinity include atomic size or effective nuclear charge, stability of electron configuration carried by the anion, and the charge carried by the anion.

78
Q

Why are second electron affinities always positive?

A

Second electron affinities are always positive because the large negative charge present on the anion creates a strong repulsive field, which repels the incoming electron. Energy is added to overcome the repulsive force, resulting in a positive overall energy change.

79
Q

What is the definition of second electron affinity?

A

Second electron affinity is the net energy absorbed or gained when one mole of an anion having a double charge is formed from one mole of a singly charged anion in the gaseous isolated state.

80
Q

Describe the first electron affinity.

A

The first electron affinity is the energy lost or gained when 1 mole of atoms in the free gaseous state gain 1 mole of electrons into the outermost shells to form 1 mole of gaseous singly charged anions

81
Q

Define electron affinity.

A

Electron affinity is the energy change that occurs when an atom gains an electron to form a negative ion.

82
Q

What causes the irregularity in ionization energy between oxygen and nitrogen?

A

The electron configuration of oxygen leads to increased repulsions and decreased binding energy, making it easier to remove electrons compared to nitrogen

83
Q

Explain why the ionization energy of boron is less than that of beryllium.

A

The electron to be removed from beryllium is in a fully filled 2s orbital, which provides extra stability compared to the incomplete 2p orbital of boron.

Year 10 Chemistry (6 decks)

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