Periodic Table Unit Flashcards

1
Q

What happens to ionization energy as it goes across a period?

A

It increases.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

Who was the first to publish the periodic table of elements?

A

Dmitri Mendeleev

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

How did he group the elements?

A

In columns by similar properties in order of increasing mass.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

How is the modern periodic table organized?

A

In the modern periodic table, elements are arranged in order of increasing atomic number. elements with similar properties are in the same column.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

What determines the chemical properties of an element?

A

The number of valence electrons.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

What are electron dot diagrams?

A

Uses the symbol of the element and dots to represent the electrons in the outer energy level.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

What are the characteristics of metals?

A

They are lustrous (shiny) solids, generally high melting points and densities, malleable (can be shaped), ductile (can be pulled into wires), conductors of heat and electricity, and they lose electrons.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

what are the characteristics of nonmetals?

A

They are poor conductors of heat and electricity, often dull and brittle, insulators, and gains electrons.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

How does distance affect the affinity for electrons?

A

The shorter the distance between the nucleus and the valence electrons, the more strongly they are held by the atom.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

What is the relationships between number of electrons and the force holding it?

A

The more electrons between the nucleus and a valence electron, the weaker the force holding it.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

What is ionization energy?

A

The energy required to remove an electron from an atom. The greater the attraction the greater the ionization energy, the stronger the attraction between the nucleus and the valence electrons, the harder it is to remove electrons.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

What are non-metals’ ionization energy/chemical properties?

A
  • They have a strong attraction between nucleus and valence electrons (high ionization energy) difficult to remove electrons
  • They have high electronegativity
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

What are metals’ ionization energy/chemical properties?

A
  • They have low ionization energy
  • They lose electrons when bonding (metal atoms nuclei have a weak attraction for their valence electrons and so lose them)
  • Alkali and alkaline earth metals have little valence electrons so they are the most reactive of the metals
  • Metals are found combined in nature
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

What happens to atomic radius as it goes down a group?

A

It increases.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

What happens to atomic radius as it goes across a period?

A

It decreases.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

What happens to electronegativity as it goes down a group?

A

It decreases.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

What happens to electronegativity as it goes across a period?

A

It increases.

18
Q

What happens to ionization energy as it goes down a group?

A

It decreases.

19
Q

What is atomic radius?

A

The distance from the nucleus to the outermost electron.

20
Q

What is electronegativity?

A

A measure of the tendency of an atom to attract a bonding pair of electrons.

21
Q

Why does atomic radius decrease across a period?

A

There are a number of protons in the nucleus so the nucleus more strongly attracts its orbital electrons and pulls them closer to it, while electron shielding remains similar.

22
Q

Why does atomic radius increase across a group?

A

Increasing number of energy levels.

23
Q

Why does electronegativity decrease down a group?

A

As you move down a group, the radius increases as more electrons shells are added. Since the outer electrons (those involved in bonding) are farther from the nucleus, they will feel the “pull” of the nucleus less. Larger atoms have lower electronegativity.

24
Q

Why does electronegativity increase across a period?

A

As we move across the period, the effective nuclear charge increases and the atomic size decreases.

25
Q

Why does ionization energy increase across a period?

A

The nuclear charge of the nucleus is increasing, and the atomic radius is decreasing. These two things increase the pull on the electrons from the nucleus, which makes it require more energy to remove an electron.

26
Q

Why does ionization energy decrease down a group?

A

Since there are new layers being added the valence electrons are farther from the nucleus and therefore the nucleus has a weaker hold on those electrons so they are easier to remove (less energy).

27
Q

What happens when you read down a group?

A

The valence electrons are further and further away from the nucleus, valence electrons are removed more easily.

28
Q

What happens to ionization energy as you move down a group?

A

It decreases.

29
Q

What is the general level of electronegativity of metals?

A

They have weak attraction for their valence electrons, so they have low electronegativity.

30
Q

What happens when metals lose their valence electrons?

A

The ion produced is smaller than the original atom.

31
Q

What happens when a nonmetal atom becomes an ion?

A

It gains electrons, so the radius increases, making it bigger than the original atom.

32
Q

What are monatomic elements?

A

Elements that only contain one atom. Noble gasses.

33
Q

What are diatomic elements?

A

H2, N2, O2, F2, Cl2, Br2, I2 - HONCLBRIF

34
Q

What are the elements phases at STP?

A

Most are solid, mercury and bromine are liquid, and the noble gases plus hydrogen oxygen nitrogen chlorine and fluoride are gases. HONCLF

35
Q

What are monatomic gases?

A

Single atoms not bonded to each other (noble gasses).

36
Q

Which element has the most metallic character and which has the most nonmetallic character?

A

Francium has the most metallic character and Fluorine has the most nonmetallic character.

37
Q

What happens to metallic character as we read down and to the left of the table?

A

It increases.

38
Q

What happens to metallic character as we read up and to the right of the periodic table?

A

It decreases.

39
Q

What happens to nonmetallic character as we read down and to the left of the table?

A

It decreases.

40
Q

What happens to nonmetallic character as we read up and to the right of the periodic table?

A

It increases.

41
Q

What is metallic character?

A

Metallic character refers to the level of reactivity of a metal. Metals tend to lose electrons in chemical reactions, as indicated by their low ionization energies.