Chapter 2 Atomic Structure And The Periodic Table Flashcards Preview

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Flashcards in Chapter 2 Atomic Structure And The Periodic Table Deck (20)
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0
Q

What are monatomic ions?

A
  • They are the charged participles that form when an atom gains or loses electrons.
  • This means the electron configurations of an atom and it’s ions must be different.
1
Q

Define isotopes.

A

Isotopes are atoms with the same number of protons (same element) but a different number of neutrons.

2
Q

Why do the atomic radii decrease from

Left to right of the periodic table?

A
  • It is primarily due to the atom’s increasing nuclear charge.
  • A higher nuclear charge (positive) increases the attraction of electrons (negative) bringing them, on average, closer to the nucleus thus resulting in a decrease in atomic radius.
3
Q

Why do the atomic radii increase down any group of a periodic table?

A
  • This is due to the outer electrons located on a higher electron shell.
  • Electrons in higher shells are on average further from the nucleus and this causes the atoms further down a group to have a larger radius.
4
Q

Define Ionisation energy.

A
  • It is a measure of how strongly the nucleus holds onto it’s electrons.
  • Importantly Ionisation energy greatly affects an element’s tendency to form positive or negative ions.
  • This in turn has implications for the type of bonding the element will undergo with itself and with other elements.
5
Q

Define First Ionisation Energy.

A

It is a measure of the minimum amount of energy required to remove the single most loosely bound electron from an atom in the neutral gaseous state.

6
Q

What 3 factors determine the trends in first Ionisation energy of the elements?

A
  1. The atom’a nuclear charge
  2. The distance between the nucleus and the outermost electron (atomic radius)
  3. Shielding by inner electrons
7
Q

Explain how the atom’a nuclear charge determines the first Ionisation energy of an element.

A
  • An atom’s nuclear charge depends upon it’s number of protons (atomic number).
  • A greater nuclear charge means electrons are attracted more strongly to the nucleus and so Ionisation energy increase with increasing nuclear charge.
8
Q

Explain how the distance between the nucleus and the outermost electron (atomic radius) determines the first Ionisation energy of an element.

A
  • As atomic radius increases so the strength of attraction between the nucleus (+) and the outer electron (-) decreases.
  • Thus Ionisation energy decreases with increasing atomic radius.
9
Q

Explain how the shielding by inner electrons determines the first Ionisation energy of an element.

A
  • Each electron in the atom’s electron cloud repels each other electron.
  • Thus the more electrons there are between the outer electron (the one to be removed) and the nucleus (to which it is attracted) the more easily the electron is removed and the lower the Ionisation energy.
10
Q

Why does the Ionisation energy increase from left to right of the periodic table?

A

This is due to the atom’s increasing nuclear charge and reducing atomic radius.

11
Q

Why for does the Ionisation energy decrease down a group?

A
  • The Ionisation energy decrease despite the atom’a increasing nuclear charge.
  • Instead the reducing Ionisation energy is a result of the outer electrons being located in progressively higher shells.

•Thus electrons are further from
the nucleus and Ionisation energy will decrease.

•The shielding effect is also greater for elements lower down in a group.

12
Q

Define electronegativity.

A

The property of electronegativity is defined as the ability of an atom in a molecule to attract electrons to itself.

13
Q

Why is it typical for elements with low Ionisation energy to have metallic properties?

A

This is because a low Ionisation energy is essential if atoms are to form the positive ions required for the metallic structure that givers metals their unique properties.

14
Q

How are covalent bonds formed?

A

At some point in each period the tendency for elements to form metallic structures (with the associated positive ions) is overcome by a tendency for elements to share electrons with one another by forming covalent bonds.

15
Q

How are covalent network structures formed?

A

The group 14 elements of highest Ionisation energy:
•carbon (a non-metal)
•silicon (a metalloid)
•and germanium (a metalloid)
-They achieve an octet in their valence shell, forming covalent bonds.

  • Each atom shares all four of its valence electrons with neighboring atoms to form four covalent bonds.
  • As a result, they form covalent network structures with physical properties of high hardness and brittleness, high melting and boiling points and semi conducting or non electrical conductivity.
16
Q

How is a diatomic molecule formed?

A

For eg, Group 17 elements all have seven valence electrons and thus form diatomic molecules with a single covalent bond (eg F2, Cl2, Br2, and I2).

17
Q

Explain why doesn’t diamond conduct electricity?

A

Diamond has covalent network bonding the electrons are highly localized in the covalent bond and are therefore not available to act as charge carriers.

18
Q

Explain why the atomic radius of Al is smaller than the atomic radius of Mg?

A

~> Al has a larger nuclear charge than Mg and since valence electrons in each element are in the same energy level (same distance from the nucleus), the electrons in the Al are pulled closer to the nucleus.

19
Q

Why is the atomic radius of calcium larger than that of magnesium?

A

~>The valence electrons in Ca are further from the nucleus than in Mg and since Ca & Mg have effectively the same core nucleus charge of +2 the Ca nucleus will be larger. [Shielding effect of the inner shell electrons]