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Chemistry > 3. Periodicity > Flashcards

3. Periodicity Flashcards

1
Q

How are elements arranged in the Periodic Table?

A

Elements are arranged by increasing atomic number and by changes in physical and chemical properties.

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2
Q

Group vs. Period

A

The group is the number of valance electrons (vertical)
The period is the number of shells (horizontal)

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3
Q

What are the different categories of elements in the Periodic Table? and where on the Periodic Table are they?

A

Alkali metals (group 1)
Alkaline earth metals (group 2)
Transition metals (group 3 - 12)
Halogens (group 17)
Noble gases (group 18)
Lanthanois (La to Lu)
Actinoids (Ac to Lr)

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4
Q

Non-metals

A

all on the right side of the periodic table (except H)

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5
Q

Non-metal physical properties

A

Mostly gaseous at room temperature
Dull
Brittle m
Poor electrical conductivity
Low densities
High ionization energies
High electronegativity

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6
Q

Non-metals chemical properties

A

Gain electrons to form anions
Oxidizing agents (therefore reduced themselves)
form covalent bonds with nonmetals
form acidic oxides

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7
Q

Metals

A

Found on the left side of the Periodic Table

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8
Q

Metals physical properties

A

Solid at room temperature (except mercury)
Metallic luster (shiny)
Malleable and ductile
High electrical conductivity
High density
Low ionization energy
Low electronegativity

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9
Q

Metals chemical properties

A

Loses electrons to form cations
Reducing agents (therefore oxidized themselves)
Form ionic bonds with nonmetals
Form basic oxides

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10
Q

What is electron shielding?

A

occurs when inner shell electrons shield valance electrons from full attraction to nucleus

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11
Q

What is the trend of electron shielding?

A

Constant across the period
Increases down a group (main level occupied increases)

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12
Q

What is the effective nuclear charge?

A

The net positive charge experienced by valance electrons
ENC = Atomic number - # of Shielding electrons

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13
Q

What is the trend for effective nuclear charge?

A

Increases across period (higher atomic number)
Constant down the group

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14
Q

What is ionization energy?

A

The minimum amount of energy required to remove a mole of electrons from a mole of gaseous atoms of a particular element.
Al(g) –> Al+(g) + e-

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15
Q

What is the trend for ionization energy?

A

Increases across period (increase in nuclear charge causes increased attraction between outer electron and nucleus, more difficult to remove)

Decreases down the group (electron being removed is from the furthest energy level from nucleus, easier to remove valance electrons as atomic radius increases down a group)

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16
Q

What is 1st ionization energy?

A

Energy required to remove one mole of electrons from one mole of gaseous atoms to produce one mole of gaseous 1+ ions (kj mol-1)

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17
Q

Exceptions to ionization energy trends

A

Mg>Al &. Be>B (electrons in P orbitals are higher in energy, further from nucleus than s orbital, require less energy to be removed)

N>O (electrons are in pairs, easier to remove because of repulsion)

18
Q

What is Electronegativity?

A

Measure of an atom’s ability to attract a bonding pair of electrons to itself.

19
Q

What are the trends for electronegativity?

A

Highest = Fluorine, Lowest = Francium

Increases across period (increase in nuclear charge, stronger attraction between nucleus and bonding pairs)

Decreases down group (atomic radius increases, bonding electron pairs are further from nucleus and there is less attraction)

20
Q

What is Atomic Radius?

A

Half the distance between the nuclei of identical atoms that are bonded together. (half the size of an atom)

21
Q

What are the trends for Atomic radii?

A

Decreases across the period (nuclear charge increases, electron shielding =, higher attraction of nucleus to outer electrons)

Increases down group (number of electron shells increases)

22
Q

What is the Ionic Radius?

A

Measure of the distance from the center of an ion to its outer electrons

23
Q

What are the trends for ionic radii?

A

Increases down a group (Due to increase in shells)

Decreases from Gp. 1 - 14 for +ions (Increasing number of protons pulling on the same number of electrons = attraction increasing)

Decreases from Gp. 14-17 for -ions

24
Q

Why is there a large increase in ionic radii from Si to P?

A

Si 4+ occupies 2 subshells while P 3- occupies 3

25
Q

Why are cations smaller than their parent atoms?

A

Ions have more protons that electrons, there is an increased attraction between the nucleus and electrons

26
Q

Why are negative ions larger than their parents ion?

A

The ion has ore electrons than protons, there is a weakened attraction between the nucleus and its electrons.

27
Q

What is electron affinity?

A

The energy change associated with the addition of an electron to a gaseous atoms
X(g) + e- -> X-(g)

28
Q

Trends of electron affinity

A

Decreases down the group

29
Q

Is second electron affinity endothermic or exothermic?

A

Endothermic (positive)
Due to the extra repulsion when adding electron to an already negative ion

30
Q

Trends down group 1 (Alkali Metals)

A

Atomic & Ionic radii increase
(more electron shells)
1st ionisation energy decreases
(valance electron is further form nucleus, therefore easier to remove)
Electronegativity decreases
(increase distance and shielding)
Melting point decreases
(atoms become larger and metallic bonds become weaker)
Reactivity increases
(valance electron is easier to lose due to shielding)

31
Q

Trend down group 7 (Halogens)

A

Atomic & Ionic radii increase
(more electron shells)
1st ionisation energy decreases
(valance electron is further form nucleus, therefore easier to remove)
Electronegativity decreases
(increase distance and shielding)
Melting point increases
(Van der Waals forces become greater with more electrons)
Reactivity decreases
(outer shell gets further from nucleus, attraction between nucleus and electrons = weaker, electron is less easily gained)

32
Q

What is the metallic character?

A

A measure of how easily an element loses a valance electron

33
Q

What are the trends for metallic character?

A

Increases down a group
Decreases across the period

34
Q

In what way do oxides change?

A

basic through amphoteric to acidic across a period

35
Q

Important equations basic oxides

A
  1. Strong base
    Na2O + H2O -> 2NaOH
  2. Weak Base
    MgO + H2O -> Mg(OH)2
    *pH will increase
36
Q

Important equations acidic oxides

A
  1. Weak acid
    P4O10 + 6H2O -> 4H3PO4
  2. Strong acid
    SO3 + H2O -> H2SO4
    *pH will decrease
37
Q

Metal + Water →

A

Metal hydroxide + hydrogen
2Li(s) + 2H2O(l) →2LiOH + H2(g)

38
Q

Metal Oxide + Water →

A

Metal Base
Li2O (s) + H2O (l) → 2LiOH

39
Q

Neutralisation Reaction

A

acid + base = salt + water

40
Q

Acid + metal →

A

Salt + Hydrogen gas

41
Q

Acid + carbonate →

A

Salt + Carbon Dioxide + Water

42
Q

Bonding properties of period 3 oxides

A

Change from ionic to covalent across period 3

Chemistry (2 decks)

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