lesson 01: the structure of the atom 01 …...chemistry 11, atoms and the periodic table, unit 07 1...

Chemistry 11, Atoms and the Periodic Table, Unit 07 1

Lesson 01: The Structure of the Atom

01 Introduction

All physical objects in the universe are composed of matter and all matter is composed of elements and all elements are composed of atoms. But what are atoms composed of?

Matter


02 Modern Basic Structure of the Atom Atoms are made up of three types of sub atomic particles…

electrons: negatively charged particles which orbit the nucleus

protons: positively charged particles which are located in the nucleus

neutrons: particles with no charge located in the nucleus

An Atom of the Element Carbon


Atoms are not all the same, since any two atoms of…

one particular element will always be identical since both would have the same number of sub atomic particles

different elements will always be different since both would have differing numbers of sub atomic particles

Two Atoms, Two Elements

Hydrogen Atom Helium Atom


03 Atomic Number, Atomic Mass, Ions and Isotopes The way in which all known elements are displayed is using the periodic table of elements…

For a neutral atom, there are three pieces of information are usually displayed on the periodic table…

element symbol

atomic number is the number of protons or electrons there are in the nucleus

atomic mass is the number of neutrons added to the number of protons


Atoms of the same element will always have the same number of…

electrons

protons

neutrons

Two Atoms of Hydrogen

H Atom H Atom

1 proton

1 electron 1 proton

1 electron

Example 01

How many electrons, protons and neutrons are in a neutral atom of C?

carbon carbonatomic number protons electrons 6

carbon carbon carbon

carbon carbon carbon

carbon

atomic mass protons neutrons

neutrons atomic mass protons

neutrons 12 6 6


Example 02

How many electrons, protons and neutrons are in a neutral atom of Fe?

iron

iron ironatomic number protons electrons 26

iron iron iron

iron iron iron

iron



neutrons 56 32 06

Atoms of the same element…

having different numbers of electrons but having the same number of protons and neutrons

are called ions

Three Forms of the Element Hydrogen

H Ion H Atom H

Ion

1 proton 1 proton

1 electron 1 proton

2 electrons

positively charged hydrogen ion

neutral hydrogen atom

negatively charged hydrogen ion


Example 03

How many electrons, protons and neutrons are in the ion

2Fe ?

iron

iron

atomic number protons

atomic number electr

26

2ns 4o 26 2

iron iron iron

iron iron iron

iron



neutrons 56 32 06

Atoms of the same element…

having the same number of electrons and protons, but differing by the number of neutrons

are called isotopes…

Three Isotopes of Hydrogen

Hydrogen Deuterium Tritium

1 proton 1 electron

1 proton 1 neutron 1 electron

1 proton 2 neutrons 1 electron

most common isotope

less common isotope

less common isotope


04 Modern Detailed Structure of the Atom

Electron Arrangement Electrons are arranged around a nucleus in shells…

In most cases (not all)…

a shell will have subshells, and

a subshell will have its own orbitals


Quantum Numbers Quantum numbers determine the location and number of electrons in a particular shell… There are four quantum numbers…

Location and Number of Electrons

Name Symbol Meaning Values

principle quantum number

n energy level 1, 2, 3…

angular momentum quantum number

l shape 0, 1, 2, 3...

s, p, d, f...

magnetic quantum number

lm orientation

s: 1 orbital

p: 3 orbitals

d: 5 orbitals

f: 7 orbitals

spin quantum number sm electron spin

1+

2 or

1-2


Orbital Shapes

Since a maximum of two electrons can exist in any atomic orbital the following results…

Number of Electrons Per Shell and Sub-Shell

Shell Sub-Shell(s) Maximum

Electrons Per Sub-Shell

Total Number of Electrons

1 s 2 2

2 s, p 2, 6 8

3 s, p, d 2, 6, 10 18

4 s, p, d, f 2, 6, 10, 14 32


Electron Configurations How do we write the exact distribution for all electrons within a particular atom? This depends on two factors…

how many electrons an atom has, and...

how many electrons each shell and sub-shell can accommodate

The Aufbau Principle is used to determine the electron configuration of an atom. According to the principle, electrons fill orbitals…

lowest energy levels first, followed by

higher levels next


To determine the electron distribution, a useful diagram is commonly used…

Orbital Filling Order for Multi-Electron Atoms


It is also useful to realize that the periodic table of the elements has been constructed with electron configuration in mind…

Orbital Filling and the Periodic Table

Knowing this, it is now possible to generate electron configurations for atoms of any element...


Example 04

Determine the electron configuration for the following atoms…

Atomic Number

Element Structure

1 Hydrogen 11s

2 Helium 21s

3 Lithium 122s1s

4 Beryllium 222s1s

5 Boron 2 2 11s 2s 2p

6 Carbon 2 2 21s 2s 2p

7 Nitrogen 2 2 31s 2s 2p

8 Oxygen 2 2 41s 2s 2p

9 Fluorine 2 2 51s 2s 2p

10 Neon 2 2 61s 2s 2p


Electron Configurations and Core Notation Displaying electron configurations for elements following a noble gas is done using core notation. Core notation is a way of displaying electron configurations in a kind of short form.

core electrons are referred to as those electrons nearest noble gas having an atomic number less than that which is being considered

outer electrons are those outside the core


Example 05

Determine the electron configuration using core notation for the following…

Atomic Number

Element Structure

10 Neon 2 2 61s 2s 2p

Neon

(core notation) [Ne]

11 Sodium 2 2 6 11s 2s 2p 3s

Sodium (core notation)

1[Ne]3s

12 Magnesium 2 2 6 21s 2s 2p 3s

Magnesium

(core notation) 2[Ne]3s

13 Aluminum 2 2 6 2 11s 2s 2p 3s 3p

Aluminum (core notation)

2 1Ne 3s 3p

14 Silicon 2 2 6 2 21s 2s 2p 3s 3p

Silicon (core notation)

2 2Ne 3s 3p


Electron Configurations and Ions Displaying electron configurations for ions is straightforward. All the needs to be done is…

determine the number of electrons in the atom that forms the ion, and then

add an electron for each negative change

remove an electron for each positive charge

Example 06

Determine the electron configuration for Br using core notation.

first determine the total number of electrons in bromine and then make sure to add an extra electron due to the presence of one negative charge

bromineatomic number electrons 35 61 3

determine the nearest noble gas

argon

place electrons in the core and remainder outside the core

2 10 5Ar 4s 3d 4p


There are two exceptions to the configurations of elements up to krypton. These two exceptions are due to the fact that there is greater stability in having either…

a filled sub-shell

exactly half filled sub-shell

Electron Configuration Exceptions

Incorrect Correct

Chromium 2 4Ar 4s 3d 1 5Ar 4s 3d

Copper 2 9Ar 4s 3d 1 10Ar 4s 3d


Valence Electrons – Accessible or Reactable Electrons Valence electrons are all accessible electrons or all electrons in outer shells even if the outer shell is completely closed (filled). Core electrons are inaccessible electrons or electrons of the nearest noble gas.

Accessible / Valence and Inaccessible / Core Electrons

element core electrons valence electrons

lithium He 12s

inaccessible accessible

Valence electrons are important in determining how the atom reacts chemically with other atoms…

complete shell: atoms with a complete or closed shell of valence electrons tend to be chemically inert

over-filled shell: atoms with one or two valence electrons more than a closed shell are highly reactive because the extra electrons are easily removed to form positive ions

under-filled shell atoms with one or two valence electrons fewer than a closed shell are also highly reactive because of a tendency either to gain the missing electrons and form negative ions


Example 07

Determine the number of valence electrons in the following atoms...

Ato

mic

Num

ber

Ele

me

nt

Str

uctu

re

Vale

nce

Ele

ctr

ons

1 Hydrogen 11s 1

2 Helium 21s 2

3 Lithium 1He 2s 1

4 Beryllium 2He 2s 2

5 Boron 2 1He 2s 2p 3

6 Carbon 2 2He 2s 2p 4

7 Nitrogen 2 3He 2s 2p 5

8 Oxygen 2 4He 2s 2p 6

9 Fluorine 2 5He 2s 2p 7

10 Neon 2 6He 2s 2p 8


Lesson 02: The Periodic Table

01 Introduction to the Periodic Table

The periodic table of the chemical elements is a method of displaying the chemical elements. The invention of the periodic table is generally credited to Russian chemist Dmitri Mendeleev in 1869. Mendeleev intended the table to illustrate recurring (periodic), trends in the properties of the elements…

Mendeleev's 1871 periodic table

http://en.wikipedia.org/wiki/File:Mendelejevs_periodiska_system_1871.png


Today, the periodic table is now common within chemistry, providing an extremely useful way to classify, systematize and compare all the many different chemicals and their behaviours. The current periodic table, as of January 2008 contains 117 elements.

Modern Periodic Table


02 Arrangement of Elements

Groups, Families or Columns A group or family is a vertical column in the periodic table. Groups are considered the most important method of classifying the elements.

Groups are numbered numerically 1 to 18 from the left to right. Some of these groups have been given trivial names, such as…

alkali metals (group 1)

alkaline earth metals (group 2)

pnictogens (group 15)

chalcogens (group 16)

halogens (group 17)

noble gases (group 18)

transition metals (group 3 through group 12)


However, other groups, are referred to simply by their group numbers, since they display fewer similarities and/or vertical trends.

Periods or Rows A period is a horizontal row in the periodic table. Although groups are the most common way of classifying elements, there are regions where horizontal trends are more significant than vertical trends.


Blocks Because of the importance of the outermost electron shell, the different regions of the periodic table are sometimes referred to as blocks, named according to the sub-shell in which the last electron resides…

s-block

p-block

d-block

f-block

Periodic Table Blocks


03 Trends in the Periodic Table

Atomic Radius

As an atom's atomic number increases, more protons are added to the nucleus, while more electrons are stacked in the shells and sub shells around the atom. The trends are…

left to right (radius decreases): electrons become more tightly held to the nucleus because they are not being added to new energy levels

up to down (radius increases): electrons become more loosely held to the nucleus because they are being added to new energy levels


Electronegativity

Electronegativity is the desire of an atom to take another atom's electrons. The trends are…

left to right (electronegativity increases): greater desire to fill missing valence shell electrons

up to down (electronegativity decreases): greater radius, increased shielding effect


Ionization Energy

The ionization energy is the energy required to remove an outermost electron. The trends are…

left to right (ionization energy increases): greater desire to fill missing valence shell electrons

up to down (ionization energy decreases): greater radius, increased shielding effect


Note Regarding the Shielding Effect The shielding effect describes the attraction between an electron and the nucleus in any atom with more than one electron shell. Shielding effect can be defined as a reduction in the effective nuclear charge on electrons due to a difference in attractive forces of between outer electrons on the nucleus.

Shielding Effect


Lesson 03: Chemical Bonding

01 Introduction

Chemical reactions involve…

making bonds, and

breaking bonds

Chemical Reactions are a Result of Chemical Bonding

Before we can understand how a chemical reaction takes place, it

is essential that we know…

what bonds are, and

how they form

Once we have discussed these two topics we also need to

discuss…

what makes bond strong or weak, and

the implications of both


02 Review

There are three terms you need to be aware of and careful not to

confuse…

valence electrons

core electrons

valence

Valence Electrons and Core Electrons

Valence electrons are those electrons in the outermost shell of an

atom. Core electrons are all electrons inside the outermost shell of

an atom.

Valence and Core Electrons Example: Carbon Atom

Valence Electrons Core Electrons

4 2

Valence electrons are important in determining two things…

IF an atom will react with other atoms, AND

HOW an atom will react with other atoms


Atoms can have a…

completely filled outer shell: tend to be chemically inert and

thus unreactive

Completely Filled Outer Shell: Helium:

He

over filled outer shell: one or two valence electrons more

than a complete shell are highly reactive because the extra

electrons are easily removed to form positive ions

Over Filled Outer Shell: Lithium:

1He 2s


under filled outer shell: one or two valence electrons less

than a complete shell are also highly reactive because of a

tendency to gain the missing electrons and form negative

ions or to share electrons and form covalent bonds

Under Filled Outer Shell: Fluorine

2 5He 2s 2p


Valence

Valence, on the other hand, is the number of electrons gained or

lost to make the outermost shell of an atom stable.

Filling or emptying the outermost shell results in increased

stability.

Atoms bond to form compounds so as to fill their outer shells and

thus increase their stability.

Note

Never assume, however, that the valence of a given atom is simply

determined by the number of electrons present in the outermost

shell, or its valence electrons. In reality it is much more

complicated. Instead, just memorize valence values as column

numbers on the periodic table.

http://chemistry.about.com/od/chemistryglossary/a/atomdefinition.htm


Valence Electrons and Valence

Here is a table showing the differences between…

valence electrons, and

valence

Ato

mic

Num

ber

Ele

men

t

Str

uct

ure

Val

ence

Ele

ctro

ns

Val

ence

1 Hydrogen 11s 1 1

2 Helium 21s 2 0

3 Lithium 1He 2s 1 1

4 Beryllium 2He 2s 2 2

5 Boron 2 1He 2s 2p 3 3

6 Carbon 2 2He 2s 2p 4 4

7 Nitrogen 2 3He 2s 2p 5 3

8 Oxygen 2 4He 2s 2p 6 2

9 Fluorine 2 5He 2s 2p 7 1

10 Neon 2 6He 2s 2p 8 0

11 Sodium 1Ne 3s

1 1


12 Magnesium 2Ne 3s

2 2

03 What is a Bond?

A chemical bond is the process…

responsible for the attractive interactions between atoms and

molecules

which confers stability to all chemical compounds

There are two main types of chemical bonds or forces...

intramolecular forces: strong bonding forces

intermolecular forces: weak bonding forces

These two types are commonly subdivided more-or-less as

follows...


04 Intramolecular Forces

There are three main types of intramolecular forces…

metallic bonds

covalent bonds

ionic bonds

Metallic Bonding

Metallic bonding is the type of bonding found in metallic

elements.

This is the electrostatic force of attraction between positively

charged ions and delocalised outer electrons.

http://www.bbc.co.uk/scotland/learning/bitesize/higher/chemistry/energy/bsp_rev1.shtml


Pure (Non-Polar) Covalent Bonding

Atoms in a covalent bond are held together by electrostatic forces

of attraction between positively charged nuclei and negatively

charged equally shared electrons.

When two atoms are bonded by a covalent bond and have…

the SAME electronegativity the electrons will be equally

shared between nuclei

For example, diatomic molecules such as…

nitrogen gas or 2N

oxygen gas or 2O

hydrogen gas or 2H


Polar Covalent Bonding

In most covalent compounds the bonding is polar covalent. This

type of covalent bond exhibits an unequal sharing of electrons.

When two atoms are bonded by a covalent bond and have…

DIFFERENT electronegativities the electrons will not be

equally shared between nuclei

In this case, the end result is one atom having a slight negative

charge and the other a slight positive charge.

http://www.bbc.co.uk/scotland/learning/bitesize/higher/chemistry/energy/bsp_rev1.shtml


Ionic Bonding

Ionic bonding is the electrostatic force of attraction between

positively and negatively charged ions.

The formation of ions is a result of a transfer of electron(s)

between atoms with a…

LARGE DIFFERENCE in electronegativities

Ionic bonding results from metals combining with non-metals. For

example, sodium combining with chlorine forms sodium chloride.


Bonding Spectrum

Covalent, polar covalent and ionic bonding can be considered as

forming a continuous bonding spectrum…

as the difference in electronegativity increases

the bonds become more polar (then finally ionic)


Electronegativity and Bond Type

Relative Strengths of Intramolecular Forces


05 Intermolecular Forces


Polarity is an important concept in understanding intermolecular

forces.

Molecules will either be…

polar, or

non-polar

Polarity is a situation that arises when a molecule develops a…

slight positive end due to fewer negative charges or more

positive charges

slight negative end due to more negative charges or less

positive charges

To determine whether a molecule is polar or not depends on

symmetry…

Polar VS Non-Polar

Polar Non-Polar

unequal sharing

of electrons between bonded

atoms

equal or unequal sharing

of electrons between bonded

atoms

asymmetrical

arrangement of bonded atoms

symmetrical

arrangement of bonded atoms

Polar Example Non-Polar Example

hydrogen fluoride, HF boron tri-fluoride, 3BF


There are five main types of intermolecular forces…

permanent dipole - permanent dipole

permanent dipole - induced dipole

instantaneous dipole - induced dipole

ion dipole

ion - induced dipole

Permanent Dipole – Permanent Dipole Force


This type of attraction take place when two or more neutral

permanently polar covalent molecules orient themselves so that

their positive and negative ends are close to each other.

One very important and unique case of the permanent dipole -

permanent dipole attraction is known as hydrogen bonding.

A hydrogen bond is a type of attractive intermolecular force that

exists between…

a hydrogen atom on one molecule, and

an electronegative atom such as nitrogen, oxygen or fluorine

on another molecule

A hydrogen bond results when the partial positive charge on the

hydrogen is attracted to a lone pair of electrons on another atom

which bears a partial negative charge.


The most common, and perhaps simplest, example of a hydrogen

bond is found between water molecules…

Permanent Dipole – Induced Dipole Force


This type of attraction takes place due to the interaction between a

permanent dipole and a non-polar molecule.

Instantaneous Dipole - Induced Dipole Force


This type of attraction is also known as…

London forces, or

London dispersion forces, or

dispersion forces

It is a relatively weak forces of attraction that exists between…

all substances, but

are of particular importance with nonpolar covalent

molecules and noble gases

In this process…

electron distribution in individual molecules suddenly

becomes asymmetrical (or distorted)

the newly formed dipoles now become attracted to one

another


The ease with which the electron cloud of an atom can be distorted

to become asymmetrical is called the molecule’s polarizability.

The greater the number of electrons an atom has, the farther they

will be from the nucleus, and the greater the chance for them to

shift positions within the molecule.

This means that larger nonpolar molecules tend to have stronger

London dispersion forces.

Ion - Dipole Forces and Ion - Induced Dipole Forces


An ion - dipole force consists of interactions between…

an ion, and

a polar molecule

They align so that the positive and negative forces are next to one

another, allowing for maximum attraction.

An ion - induced dipole force consists of interactions between…


an ion, and

a non-polar molecule

Like a dipole-induced dipole force, the charge of the ion causes a

distortion of the electron cloud on the non-polar molecule.

Relative Strengths of Intermolecular Forces


06 Structures Caused by Bonding


Metallic Structure

A metallic structure consists of a giant lattice of positively charged

ions and delocalised outer electrons.

Covalent Molecular Structure

A covalent molecular structure consists of discrete molecules held

together by weak intermolecular forces, and hydrogen bonds.

Network Covalent Structure


A network covalent solid is a chemical compound in which the

atoms are bonded by covalent bonds in a continuous network. In a

network solid there are no individual molecules and the entire

crystal may be considered a macromolecule.

Ionic Structure

An ionic structure consists of a giant lattice of oppositely charged

ions.


Monatomic Structure

A monatomic structure consists of discrete atoms held together by

intermolecular forces, such as the Noble gases.


07 Why Make a Bond?

In nature we find that…

Noble gases: are never found bonded to other atoms,

whereas

most other elements: are only found bonded to other

elements

The reason for this is due to electron configurations…

Noble gases: the valence shell is completely full, it cannot

accept another electron into the shell

most other elements: the valence shell is not completely full,

it can accept another electron into the shell

In essence, all atoms with unfilled valence shells…

react with other atoms to form bonds, in an attempt

to achieve a closed shell electron configuration

Example 01: Lithium Fluoride

For example, when a lithium atom and a fluorine atom meet,

lithium can achieve a noble gas configuration by donating an

electron to fluorine...


Example 02: Diatomic Fluorine

You may have asked yourself why two fluorine atoms don't come

together to perform the following reaction...

Even though the reaction may appear to be favourable because of

the production of a closed shell species, there is a better way to

achieve a Noble gas configuration… by sharing electrons.

See the next part of this lesson…


08 Lewis Structures and Covalent Bonding

A covalent bond represents a shared electron pair between

atoms…

Methane Covalent Bond

Before we can determine how an atom bonds covalently with

other atoms we need a simple method for displaying valence

electrons…


Lewis Structures Basics

Since only valence electrons are involved in bonding, all other

electrons can be ignored when it comes to determining the number

of covalent bonds between atoms. The way involves drawing

Lewis Structures.

Lewis structures are drawn using the following steps…

determine the number of valence electrons

draw one dot next to the symbol for the atom for each

valence electron

place two electrons (max) on each of the four sides of the

atomic symbol since atoms strive to achieve a full octet of

electrons

Lewis Structures for First 18 Elements

Lewis structures allow us to simplify complex multi-electron

diagrams.


Lewis Structures and Covalent Bonding (Atoms)

Lewis structures are going to display two features…

bonding (pair) electrons: drawn using a line between the two

atoms

non-bonding electrons: drawn using dots to represent each

non-bonding electron

Example 03

Take for instance the bonding that takes place between hydrogen

and bromine…

Once bonding has completed, both atoms have completely filled

outer shells…

hydrogen has two electrons

bromine has eight electrons


Example 04

The deadly gas carbon monoxide provides an interesting example

of how to correctly draw Lewis structures…

Since carbon has four electrons and oxygen has six…

using a single bond: if only one bond were to be formed

between carbon and oxygen, carbon would have five

electrons and oxygen 7

using a double bond: if two bonds were to be formed

between carbon and oxygen, carbon would have six


using a triple bond: if three bonds were to be formed

between carbon and oxygen, carbon would have eight


Such multiple bonds must be employed to explain the bonding in

many molecules. However, only single, double, and triple bonds

are commonly encountered.


Lewis Structures and Covalent Bonding (Ions)

For more complex molecules and molecular ions, it becomes

important to keep an accurate count of the number of electrons in

the molecule.

Example 05

For example, let us make a Lewis structure for 2NO .

We have…

5 electrons from nitrogen

12 from the oxygens

1 extra electron due to the negative charge

Therefore, 2NO has a total of 18 electrons and we should draw

the following Lewis structure…

If we had tried to draw the above structure without taking the

charge of the ion into account, we could not have produced a full

octet around at least one atom.


Example 05

If the 2NO ion had been positively charged, as in 2NO , we

would count the electrons as follows…

5 from nitrogen

12 from oxygens

-1 one due to the charge

Therefore 2NO has 16 electrons and we should draw the

following structure...


Lesson 04: Properties of Major Chemical Families

Nobel Gases

Atomic

Number

Element Electrons /

Shell

Reactivity MP / BP

2 helium 2

low

low

10 neon 2, 8

18 argon 2, 8, 8

36 krypton 2, 8, 18, 8

54 xenon 2, 8, 18, 18,

8

86 radon 2, 8, 18, 32,

18, 8 high

Tendency: stable.

The noble gases are a group of chemical elements with very

similar properties: under standard conditions, they are all odorless,

colorless, monatomic gases, with very low chemical reactivity.

The six noble gases that occur naturally are helium (He), neon

(Ne), argon (Ar), krypton (Kr), xenon (Xe), and the radioactive

radon (Rn).

The properties of the noble gases can be well explained by modern

theories of atomic structure: their outer shell of valence electrons

is considered to be full, giving them little tendency to participate

in chemical reactions.

Like other groups, the members of this family show patterns in its

electron configuration, especially the outermost shells resulting in

trends in chemical behavior...


The noble gases have full valence electron shells. Valence

electrons are the outermost electrons of an atom and are normally

the only electrons that participate in chemical bonding. Atoms

with full valence electron shells are extremely stable and therefore

do not tend to form chemical bonds and have little tendency to

gain or lose electrons.


Alkali Metals

Atomic

Number

Element Electrons /

Shell

Reactivity MP / BP

1 hydrogen 1 low high

3 lithium 2, 1

11 sodium 2, 8, 1

19 potassium 2, 8, 8, 1

37 rubidium 2, 8, 18, 8, 1

55 cesium 2, 8, 18, 18,

8, 1

87 francium 2, 8, 18, 32,

18, 8, 1 high low

Tendency: give up 1 electron.

The alkali metals include lithium (Li), sodium (Na), potassium

(K), rubidium (Rb), cesium (Cs) and francium (Fr). Hydrogen (H),

although a member, very rarely exhibits behaviour comparable to

the alkali metals. This group lies in the s-block of the periodic

table, which means that all its elements have their outermost

electron in an s-orbital.


electronic configuration, especially the outermost shells, resulting

in trends in chemical behaviour...

The alkali metals are all highly reactive and are never found in

elemental forms in nature. Because of this, they are usually stored

in mineral oil or kerosene.


Chemically, all of the alkali metals react aggressively with the

halogens to form ionic salts. They all react with water to form

strongly alkaline hydroxides. The vigor of reaction increases down

the group. All of the atoms of alkali metals have one electron in

their valence shells, hence their only way for achieving the

equivalent of filled outermost electron shells is to give up one

electron to an element with high electronegativity, and hence to

become singly charged cations.


Alkaline Earth Metals

Atomic

Number

Element Electrons /

Shell

Reactivity MP / BP

4 Beryllium 2, 2 low high

12 Magnesium 2, 8, 2

20 Calcium 2, 8, 8, 2

38 Strontium 2, 8, 18, 8, 2

56 Barium 2, 8, 18, 18,

8, 2

88 Radium 2, 8, 18, 32,

18, 8, 2 high low

Tendency: give up 2 electrons.

The alkaline earth metals contain beryllium (Be), magnesium

(Mg), calcium (Ca), strontium (Sr), barium (Ba) and radium (Ra).

The group lies in the s-block of the periodic table.

This specific group in the periodic table owes its name to their

oxides that give basic solutions. These oxides melt at such high

temperature that they remain solids in fires.


electron configuration, especially the outermost shells resulting in

trends in chemical behavior...

All the alkaline earth metals have two electrons in their valence

shell, so the energetically preferred state of achieving a filled

electron shell is to lose two electrons to form doubly charged

positive ions.


The alkaline earth metals have much higher melting points and

boiling points compared to the alkali metals. The alkali metals are

also softer and are more lightweight whereas the alkaline earth

metals are much harder and denser.

The second valence electron is very important when it comes to

comparing chemical properties of the alkaline earth and the alkali

metals. The second valence electron is in the same “sublevel” as

the first valence electron. This means that the elements of the

alkali earth metals group contain a smaller atomic radius and

much higher ionization energy than the alkali metals group. Even

though the alkali earth metals group contains much higher

ionization energy, they still form an ionic compound with 2+

cations. Beryllium, however, behaves differently. This is because

in order to remove two electrons from this particular atom, it

requires significantly more energy. It never forms Be2+ and its

bonds are polar covalent.


Halogens

The halogens or halogen elements are a series of nonmetal

elements comprising fluorine (F), chlorine (Cl), bromine (Br),

iodine (I), and astatine (At).

Like other groups, the candidates of this family show patterns in

its electron configuration, especially the outermost shells resulting

in trends in chemical behavior...

Atomic

Number

Element Electrons /

Shell

Reactivity MP / BP

9 fluorine 2, 7 high low

17 chlorine 2, 8, 7

35 bromine 2, 8, 18, 7

53 iodine 2, 8, 18, 18,

7

85 astatine 2, 8, 18, 32,

18, 7 low high

Tendency: gain 1 electron.

Halogens are highly reactive, and as such can be harmful or lethal

to biological organisms in sufficient quantities. This high

reactivity is due to the atoms being highly electronegative due to

their high effective nuclear charge. They can gain an electron by

reacting with atoms of other elements. Fluorine is one of the most

reactive elements in existence, attacking otherwise inert materials

such as glass, and forming compounds with the heavier noble

gases. It is a corrosive and highly toxic gas. The reactivity of

fluorine is such that if used or stored in laboratory glassware, it


can react with glass in the presence of small amounts of water to

form silicon tetrafluoride. Thus fluorine must be handled with

substances such as Teflon (which is itself an organofluorine

compound), extremely dry glass, or metals such as copper or steel

which form a protective layer of fluoride on their surface.

The high reactivity of fluorine means that once it does react with

something, it bonds with it so strongly that the resulting molecule

is very inert and non-reactive to anything else. For example,

Teflon is fluorine bonded with carbon.

Both chlorine and bromine are used as disinfectants for drinking

water, swimming pools, fresh wounds, spas, dishes, and surfaces.

They kill bacteria and other potentially harmful microorganisms

through a process known as sterilization. Their reactivity is also

put to use in bleaching. Sodium hypochlorite, which is produced

from chlorine, is the active ingredient of most fabric bleaches and

chlorine-derived bleaches are used in the production of some

paper products. Chlorine also reacts with sodium to create sodium

chloride, which is another name for table salt.

lesson 01: the structure of the atom 01 …...chemistry 11, atoms and the periodic table, unit 07 1...

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