1. Atoms and the Periodic Table
Chapter 2
Atoms and the Periodic Table

2. Classification of Matter

3. Elements
A pure substance that cannot be broken down into simpler substances by a chemical reaction
Identified by a one- or two-letter symbol
Arranged in the periodic table
Its location on the periodic table indicates a lot about its chemical properties
Can be a metal, nonmetal, or a metalloid

4. Elements Some of the element symbols will be familiar to you already
It will take time and practice to learn others

6. Periodic Table

7. Elements
Metals- located to the left of the stair step line (top located near Boron)
Shiny
Good conductors of heat and electricity
Solid at room temperature except for mercury

8. Elements
Nonmetals – except for hydrogen (H) these are located to the right of the stair step line (top located near Boron)
Do not have a shiny appearance
Generally poor conductors of heat and electricity
Usually a solid or gas at room temperature (exception: Br2 is a liquid at room temperature)

9. Elements
Metalloids – 7 elements located along the stair step line (top located near Boron)
Properties in between metals and nonmetals
Metallic shine but brittle
Boron (B)
Silicon (Si)
Germanium (Ge)
Arsenic (As)
Antimony (Sb)
Tellurium (Te)
astatine (At)

10. Classification of Matter

11. Compounds
A pure substance formed by chemically combining two or more elements
A chemical formula for a compound consists of:
Element symbols to show the identity of the elements that form a compound
Subscripts to show the ratio of atoms in the compound
H2O
H2O
C3H8
C3H8
2 H atoms
1 O atom
3 C atoms
8 H atoms

12. Compounds
Can be drawn many ways but we are always making a simplified representation of reality

13. 4 Common ways to draw compounds
H H
Line
Lewis Dot Diagram

14. Atoms All mater is composed of building blocks called atoms
Atoms are composed of three subatomic particles
Proton
Neutron
Electron

15. Atoms
The mass of an atom is very small since each proton and neutron has a mass of X
Chemists use the atomic mass unit (amu) to measure the mass of an element

16. Subatomic Particles
The nucleus is the dense core that contains the protons and neutrons
Most of the mass of an atom is in the nucleus
Electron cloud is composed of electrons that are in almost empty space around nucleus

17. Subatomic Particles Opposite charges attract each other
Like charges repel each other
Protons and electrons attract each other but electrons repel each other

18. Atomic Number
The number of protons in the nucleus of an atom is its atomic number
Every atom of a given element has the same number of protons in the nucleus
We can symbolize the number of protons with the generic letter Z
A neutral atom has no overall chare so:
Z= number of protons = electrons
Example: Lithium
Atomic number (Z) is
the number of protons
in the nucleus. 3 Li

19. Isotopes
Atoms of the same element that have a different number of neutrons
The number of protons remains constant
Different forms of the same element
The mass number (A) represents the number of protons plus the number of neutrons
A = Z + number of neutrons
Example – two forms of carbon:
C-12 has 6 neutrons
C-14 has 8 neutrons
Z = 6 for both forms of carbon

20. Example Cl Mass number (A) Atomic number (Z) # of protons = 1735 Cl
Atomic number (Z) 17
# of protons = 17
# of electrons = 17
# of neutrons = A – Z = 35 – 17 = 18

21. Atomic Weight
The weighted average of the masses of the naturally occurring isotopes of a particular element reported in atomic mass units (amu)
This information can be found on the periodic table 6 C
12.01
atomic number (Z)
element symbol
atomic weight (amu)

22. Atomic Weight Calculation 1
The weighted average of the ass of the naturally occurring isotopes
Example
What is the atomic weight of chlorine?
List each isotope, it’s mass in atomic
mass units, and it’s abundance in nature.
Step [1]
Isotope
Mass (amu)
Isotopic Abundance
Cl-35
34.97
75.78% =
Cl-37
36.97
24.22% =

23. Atomic Weight Calculation 2
Multiply the isotopic abundance by the mass
of each isotope, and add up the products.
Step [2]
The sum is the atomic weight of the element. x =
amu x =
amu
35.45 amu
amu =
4 sig. figs.
Answer
4 sig. figs.

24. Organizing the elements
Periodic table was put together over time as the elements were discovered, isolated, and studied.
Many versions exist but the most common is based on the one developed by Dmitri Mendeleev in 1869
A row in the periodic table is called a period
Elements in the same row are similar in size
A column in the periodic table is called a group
Elements in the same group have similar electronic and chemical properties

25. Periodic Table

26. Periodic Table - Groups
Alkali Metals
Alkaline Earth Metals
Transition Metals
Lanthanide & Actinide
Halogens
Nobel Gases
Very reactive
Metals except for H
+1 ions
React with Oxygen to form compounds that dissolve into alkaline solutions in water
Reactive
+2 ions
Oxygen compounds are strongly alkaline
Many are not water soluble
Metals
Form ions with several different charges (oxidation states)
Tend to form +2 and +3 ions
Lanthanides 58 – 71
Actinides
90 – 103
Actinides are radioactive
Form diatomic molecules in elemental state
-1 ions
Salts with alkali metals
Inert
Heavier elements have limited reactivity
Do not form ions
Monoatomic gases

27. Carbon is Special
Carbon’s ability to join with itself and other elements gives it a versatility not seen with any other element in the periodic table
Elemental forms of carbon include the following carbon-only structures:
diamond
graphite
buckminsterfullerene

28. Carbon
“the party animal of the atomic world, latching onto many other atoms (including itself) and holding tight, forming molecular conga lines of hearty robustness-the very trick of nature necessary to build proteins and DNA”
Bill Bryson from A Short History of Nearly Everything

29. Electronic Structure
The chemical properties of an element are determined by the number of electrons in an atom
Electrons do not move freely in space – restricted to a region with a particular energy
Electrons occupy discrete energy levels that are restricted to specific values – the energy is “quantized”
Electrons are confined to regions called the principal energy levels or shells

30. Electron Shells
The shells are numbers, n=1, 2, 3, 4… Moving out from the nucleus
Electrons closer to the nucleus are held more tightly are lower in energy
Electrons farther from the nucleus are held less tightly and are higher in energy
The farther a shell is from the nucleus, the larger its volume, and the more electrons it can hold

31. Electron Shells
Nucleus

32. Orbitals Each shell is divided into subshells made up of orbitals
These are identified as s, p, d, f
Each orbital can hold two electrons

33. Orbitals

34. Orbitals
The maximum number of electrons that can occupy a shell is determined by the number of orbitals in a shell

35. Electron Configuration
The electron configuration shows how the electrons are arranged in an atom’s orbitals
The ground state is the lowest energy arrangement
The outtermost shell is the valance shell
The electrons in the valence shell are called valence electrons
The chemical properties of an element depend on the number of electrons in the valence shell

36. Electron Configuration
Rule 1:
Electrons are placed in the lowest energy orbital beginning with the 1s orbital.
Orbitals are then filled in order of increasing energy.
Rule 2:
Each orbital holds a maximum of 2 electrons
Rule 3:
1 electron is added to each orbital until all of orbitals are half-filled
Then, the orbitals can be completely filled

37. Electron Configuration

38. Example What would the electron configuration for fluorine (F) be?
Atomic Number = 9 so it has nine electrons
Shells fill: 1s22s22p5
How many valence electrons does it have?
There are 7 electrons in the 2nd shell so there are 7 valence electrons

39. Examples Be Cl 1s22s2 1s22s22p63s23p5 valence shell: n = 2
Atomic number: 4
Atomic number: 17
1s22s2
1s22s22p63s23p5
valence shell: n = 2
valence shell: n = 3
# of
valence electrons = 2
# of
valence electrons = 7

40. Valence Electrons
Elements in the same group have similar electron configurations.
Elements in the same group have the same number of valence electrons.
The group number, 1A–8A, equals the number of valence electrons for the main group elements.
The exception is He, which has only 2 valence electrons.
The chemical properties of a group are therefore very similar.

41. Valence Electrons

42. Electron Dot Symbols
Dots representing valence electrons are placed on the four sides of an element symbol
Each dot represents one valence electron
For 1 to 4 valence electrons, single dots are used
With more than 4 valence electrons the dots are paired
Element:
# of Valence electrons:
Electron-dot symbol: H 1 C 4 O 6 Cl 7

43. Periodic Trends The size of atoms increases down a column, as the
valence e− are
farther from the
nucleus.
Increases
Decreases
The size of atoms decreases across a row, as the number of protons in the nucleus increases, pulling the valence electrons in closer.

44. Periodic Trends The ionization energy is the energy needed to remove
Decreases
The ionization energy is the energy needed to remove
an electron from a neutral atom.
Na energy Na e–
Ionization energies
decrease down a
column as the
valence e− get
farther away from
the positively charged nucleus.
Increases
Ionization energies increase across a row as the
number of protons in the nucleus increases.