Unit 5 (Periodic Trends) Flashcards
Explain these people and what they did
1. Mendeleev (1869)
2. Moseley (1913) + periodic law
Mendeleev: organized the elements by increasing atomic mass AND so that elements in the same row have similar properties
Moseley: rearranged the elements by increasing atomic number, how the modern periodic table is arranged today
Periodic law: when elements are arranged in order of increasing atomic number, there is a periodic pattern in their physical and chemical properties
Difference? Compare and define
1. periods (series)
2. Groups (Families)
Period: the horizontal rows of the periodic table
* Elements in the same period do not have the same properties, but do have the same number of occupied energy levels
* periods are numbered 1-7
Group: the vertical columns on the periodic table
* Elements in the same group have similar properties
* groups are numbered 1-18
Explain the blocks on the periodic table and what they are called
1. s and p blocks
2. d block
3. f block
s and p blocks = representative elements 2 and 6 groups long
d block = transition metals 10 groups long
f block = inner transition metals 14 groups long
Compare metals and non metals
1. staircase position
2. lusture
3. Shapeability
4. conduction/insulation
Metals
* left of staircase
* lustrous (shiny)
* malleable (hammer in thin sheets) and ductile (thin wires)
* Good heat and electricity conductors
Nonmetals
* Right of staircase including hydrogen
* Non-lustrous (dull)
* Brittle (breaks easily)
* Poor conductors and good insulators
Explain metalloids and what they do.
1. Definition
2. Conductivity
Definiton touch the staircase and have properties of both metals and non metals. (Note Al is a metal, there’s 6 in total)
They are semi conductors. They noramlly do not conduct electricity but will conduct at high temperatures or when certain substances are added
Explain the concept of atom size and radius
1. if atoms have a fixed radius
2. how the radius of an atom can only be found
3. what this is called
Atoms don’t have a fixed radius.
The radius of an atom can only be found by measuring the distance between the nuclei of two touching atoms, and then halving that distance
1/2 of the distance between the nuclei of two like atoms is called Atomic Radius
What is the radius of an atom dependent on?
* the 2 things
nuclear charge (magnitude of attraction from nucleus).
Distance between electrons and nucleus
Atomic radius
What the periodic trend and group trend for atomic radius.
1. periodic trend and WHY it changes the way it does
2. group trend and WHY it changes the way it does
Period trend: (Across) atomic size decreases from left to right
* These atoms have the same number of occupied energy levels, but a greater number of protons (increased nuclear charge)
* The energy levels are more attracted to the nucleus and are ”pulled in”
Group trend: (Down) atomic size will increase as you move down a group
* Moving down a group there is a greater number of occupied energy levels (greater distance)
Explain the shielding effect
* What it is
* What it involves and how it affects atomic radius
The inner electrons shield or black the attraction from the nucleus
While there is an increase in number of protons down a group (nuclear charge) this does not affect the magnitude of attraction from the nucleus due to this effect
Define ionization energy and give an example.
Also explain the 2 things that this depends on
the amount of energy required to remove an electron from a gaseous atom.
Example: Li (g) -> Li ^+1 (g) + e-
depends on
* Distance between valence electron and nucleus
* Effective nuclear charge (# of protons)
Ionization energy
Explain the group trend (Down) and periodic trend (across) for I.E.
1. group trend (what happens and why)
2. periodic trend (what happens and why)
Group trend
* I.E. decreases down a group due to the valence electrons being further from the nucleus.
* More difficult to remove e- = e- closer to nucleus = higher I.E. Easier to remove e- = e- further away from nucleus = lower I.E.
Periodic trend
* I.E. increases as you move across a period due to an increased effective nuclear charge
* Electrons are more attracted to the nucleus due to an increased number of protons and therefore more I.E. is required to remove the electron
Explain how it requires more energy to remove and 2nd or 3rd electron and what those things are
1. explanation of the general thing
2. 1st I.E., 2nd I.E., explanation
3. expression to help remember
It always requires more energy to remove a - 2nd or 3rd electron because they are more attracted to a positive ion than to a neutral atom.
- 1st I.E. – Energy required to remove the 1st electron.
- 2nd I.E. - Energy required to remove the 2nd electron.
- 3rd I.E. - Energy required to remove the 3rd electron.
1st I.E. < 2nd I.E. < 3rd I.E.
Explain the 1st, 2nd, and 3rd I.E.s for sodium and magnesium
I don’t know if we need this but idk
Sodium
1. 495.9 kJ
2. 4560 kJ
Magnesium
1. 738.1 kJ
2. 1450 kJ
3. 7730 kJ
When is there:
1. a very large increase in ionization energy
2. why is this the case
There is a very large increase of ionization energy whenever an electron is removed from an atom/ion that is isoelectronic with a noble gas
This is due to the fact that the electron being removed is an inner (core) electron
Explain Cations
1. are they larger or smaller than a neutral atom when made?
2. why is the above the case??
3. what happens to the size?
always smaller
as electrons are lost, sometimes there is a loss of an energy level. (Decreased distance results in more attraction.) There will be less electron repulsions and therefore greater attraction to the nucleus.
The more electrons lost, the smaller the ion becomes
* Mg > mg^ +1 > Mg^+2 (smallest in size)
