Periodic Table/Electron Configuration/Periodic trend Quiz

1. What’s the electron configuration of neon?

2. What atomic model is the electron configuration based on?

3. What’s the name of the group that neon is in?

4. List the elements that have both metal and nonmetal characteristics. What do you call these elements?

5. Between S and Cs, which element has bigger radius?

Why learn electron configuration?

• Let’s think about what we learned in matter…• What is physical property?

• What is chemical property?

• Property is determined by valence e-• How reactive is element?

• How do we figure out how many valence electrons are in element?

• ELECTRON CONFIGURATION!!

WHAT DETERMINES HOW REACTIVE AN ELEMENT IS?

Valence electron

• Def) electrons in the outermost electron shell of an element.

• There are 3 ways to figure it out!

• Let’s do quick review on electron configuration!• Carbon’s # of e-=

• E- configuration=

• Let’s look the Periodic Table

Octet rule

• Elements would like to have 8 electrons in their outermost energy level (highest n value). This means that elements try to get a s2p6 configuration in their highest n value energy level.

• Which group has all the s & p filled with e-?

• Many transition metals also lose some of their “d” electrons in addition to the electrons needed to satisfy the octet rule

One way to divide periodic table is by using electron

configuration

Here’s trivia question! Who came up with basics of periodic table?

• Dmitri Mendeleev!! Pg. 390

Noble Gases

• Outermost s and p sublevels are filled

• Pg.395

• Group 18

• Also called as inert gas because of how it isn’t very reactive

Representative elements

• Outermost s or p sublevel is only partially filled

• Which groups are representative elements?

• lithium, sodium, potassium, rubidium, cesium, and francium

• Alkali metals are so named because they are metals that react with water to make alkaline solutions.

• Because the alkali metals have a single valence electron, they are very reactive.

Group 1A)alkali metals

Group2 Alkaline-Earth metals

• Which element are alkaline-earth metals?

• What’s the configuration?

• Slightly less reactive than alkali metals

• Usually found in compound

• Takes more energy than alkali metals to fulfill octet rule

Gp17 Halogens

• Most reactive group of nonmetal elements

• Why is it so reactive?

• Because the alkali metals have one valence electron, they are ideally suited to react with the halogens

• Salt former

Transition metals

• Outermost s sublevel and nearby d sublevels contain electrons.

• Group B elements

• Can use inner shell before using the outer shell to bond

• Can lose 1-3 valence electrons depending on the element with which it reacts with

The Inner transition metal

• Outermost s sublevel and nearby f sublevel generally contain electrons

• many are man made

Another way to divide them is by its property

Metals, Nonmetals, Metalloids

Metals, Nonmetals, Metalloids

• There is a zig-zag or staircase line that divides the table.

• Metals are on the left of the line, in blue.

• Nonmetals are on the right of the line, in orange.

Metals, Nonmetals, Metalloids

• Elements that border the stair case, shown in purple are the metalloids or semi-metals.

• There is one important exception.

• Aluminum is more metallic than not.

Metals

• Metals are lustrous (shiny), malleable, ductile, and are good conductors of heat and electricity.

• They are mostly solids at room temp.

Nonmetals

• Nonmetals are the opposite.

• They are dull, brittle, nonconductors (insulators).

• Some are solid, but many are gases, and Bromine is a liquid.

Metalloids• Metalloids, aka semi-metals

are just that.

• They have characteristics of both metals and nonmetals.

• They are shiny but brittle.

• And they are semiconductors.

• What is our most important semiconductor?

Periodic Trends

• The first and most important is atomic radius.

• Using your knowledge of mathematics, figure out how you could measure the radius of an atom.

• Radius is the distance from the center of the nucleus to the “edge” of the electron cloud.

• What is this statement assuming?

Atomic Radius

• Since a cloud’s edge is difficult to define, scientists use define covalent radius, or half the distance between the nuclei of 2 bonded atoms.

2.86 Å

Atomic Radius

• The trend for atomic radius in a vertical column is to go from smaller at the top to larger at the bottom of the family.

• Based on what you know, how would you explain the growing size of atom?

• With each step down the family, we add an entirely new energy level to the electron cloud, making the atoms larger with each step.

Atomic Radius

• The trend across a horizontal period is less obvious.

• What happens to atom as we step from left to right?

• Each step adds a proton and an electron.• What does this mean in terms of charge?

Atomic Radius

• The effect is that the more positive nucleus has a greater pull on the electron cloud.

• The increased attraction pulls the cloud in, making atoms smaller as we move from left to right across a period.

Effective Nuclear Charge

• What keeps electrons from simply flying off into space?

• Effective nuclear charge is the pull that an electron “feels” from the nucleus.

• The closer an electron is to the nucleus, the more pull it feels.

• As effective nuclear charge increases, the electron cloud is pulled in tighter.

Atomic Radius

• On your periodic table sheet, draw arrows like this:

Let’s practice!!!

• Which of the following elements has the largest atomic radius?

1. carbon (C) 2. silicon (Si) 3. germanium (Ge) 4. tin (Sn)

• Which of the following elements has the largest atomic radius?

1. nitrogen 2. oxygen 3. fluorine 4. neon

• The atomic radius of F, Br, and I are 64, 114, and 138 pm respectively. From this information (and not your book) estimate a reasonable atomic radius of Cl.

1. 53pm 2. 89pm 3. 126pm 4.162pm

Ions

• When an atom gains an electron, it becomes negatively charged (more electrons than protons ) and is called an anion.

• In the same way that nonmetal atoms can gain electrons, metal atoms can lose electrons.

• They become positively charged cations.

Ions

• Here is a simple way to remember which is the cation and which the anion:

This is a cat-ion.This is Ann Ion.

He’s a “plussy” cat!She’s unhappy and negative.

+ +

Ionic Radius• Cations are always smaller than the original atom.

• Who can explain why cation is smaller than original atom?

• The entire outer PEL is removed during ionization.

• Conversely, anions are always larger than the original atom.

• Electrons are added to the outer PEL.

Cation Formation

11p+

What element is this? How may Valence e-?

Valence e- lost in ion formation

Effective nuclear charge on remaining electrons increases.

Remaining e- are pulled in closer to the nucleus. Ionic size decreases.

Result: a smaller sodium cation, Na+

Anion Formation

17p+

Chlorine atom with 7 valence e-

One e- is added to the outer shell.

Effective nuclear charge is reduced and the e- cloud expands.

A chloride ion is produced. It is larger than the original atom.

Let’s practice

• Which has the larger radius?

• 1. Na2. Na+

• Which has the larger radius?

• 1. O 2. O2-

• Which has the largest radius?

• 1. K 2. K+ 3. Rb 4. Rb+

Shielding

• As more PELs are added to atoms, the inner layers of electrons shield the outer electrons from the nucleus.

• The effective nuclear charge (enc) on those outer electrons is less, and so the outer electrons are less tightly held.

Ionization Energy• If an electron is given enough energy (in the

form of a photon) to overcome the effective nuclear charge holding the electron in the cloud, it can leave the atom completely.

• The atom has been “ionized” or charged.

• The number of protons and electrons is no longer equal.

Ionization Energy• The energy required to remove an electron

from an atom is ionization energy. (measured in kilojoules, kJ)

• The larger the atom is, the easier its electrons are to remove.

• Why?• Ionization energy and atomic radius are

inversely proportional.

Ionization Energy (Potential)

• Draw arrows on your help sheet like this:

Let’s practice

• Which element has greater ionization energy?

• 1. N 2. O

• From which element is it harder to pull electron away?

• 1. Be 2. B

• Which one has greatest ionization energy?

• 1. beryllium (Be) 2. strontium (Sr) 3. calcium (Ca) 4. magnesium (Mg)

• Which one has strongest shielding effect?

• 1. beryllium (Be) 2. strontium (Sr) 3. calcium (Ca) 4. magnesium (Mg)

Electronegativity

• Electronegativityis the tendency for atom to attract electron when they are chemically bonded.(also measured in kJ).

EN depends on…

• Atomic size• Greater the size of the atom, the lesser the

electronegativity

• Nuclear charge• The greater the nuclear charge, the greater the

electronegativity

Electronegativity

• Draw arrows on your help sheet like this:

Let’s practice!

• Which has the highest electronegativity?

• 1. Na 2. Al 3. S 4. Cl

• Which element attracts electron least?

• 1. F 2. I 3. Br 4. Cl