1. Chapter 4 Elements, Atoms & Ions1

2. 4.1 Elements
4.1 The Elements
Substances that cannot be decomposed into simpler substances by chemical or physical means.
Each element consists of atoms having the same atomic number.
Over 112 known. Of first 92 (H through U), 90 are found in nature; rest are man-made.
Abundance is the percentage found in nature
Oxygen is most abundant element (by mass) in the crust of the earth and in the human body (Tables 4.1, 4.2).
Iron, however, is the most abundant element (by mass) throughout the entire earth. (About 35% of Earth’s mass) 2

3. 4.3 Dalton’s Atomic Theory
4.2 Symbols for Elements
Each element has a unique symbol
The symbol of an element may be one letter or two
If two letters, the second is lower case
4.3 Dalton’s Atomic Theory
The ancient Greeks (Democritus, c. 400 BCE) were first to develop the idea of the atom as the smallest possible particle of matter—something that could not be subdivided any further.
Did not know about parts of an atom
This idea of the atom as an indivisible particle of matter persisted up until the 19th century . 6

4. Based on his observations and experiments, Dalton proposed that:
John Dalton, an English scientist, in the early 1800s, developed the first atomic theory.
Based on his observations and experiments, Dalton proposed that:
Elements are made of tiny particles called atoms.
All atoms of a given element are the same
The atoms of any given element are different from those of any other element.
Atoms of different elements can combine to form compounds. A given compound always contains the same relative number and type of atoms.
Atoms are indivisible in chemical processes. They cannot be created or destroyed. They can only be changed in how they are grouped together.
Law of Constant Composition: any given compound (a pure substance) has the same composition, regardless of where it is from.
Ex:
A compound is a distinct substance composed of atoms of two or more elements.
Compounds always contain the same relative numbers of atoms of each element
Compounds always contain the same relative masses of each element
A formula describes a compound by indicating the number and type of each atom present in the simplest unit of the compound 6

5. 4.4 Formulas of Compounds Rules for writing formulas (pg 90)
each element is represented by its letter symbol
the number of atoms of each element is written to the right of the element as a subscript
When only one atom of a type is given, the subscript 1 is not written
polyatomic groups are placed in parentheses (if more than one)
Examples:

6. 4.5 The Structure of the Atom
J. J. Thomson
Discovered the electron showing that the atoms has parts
Cathode ray tube experiment
“Plum Pudding” Model of the Atom

7. https://youtu.be/kBgIMRV895w
Ernest Rutherford
-- In 1910, his gold foil experiment showed that atoms do not have the same consistency throughout.
-- All the mass is located in the center (a positively-charged nucleus) with the electrons outside the nucleus.

8. 4.6 The Modern Concept of Atomic Structure
Atoms are the fundamental units of which elements are composed. The smallest unit of matter that retains all properties of an element.
Atoms are composed of three main parts: protons, neutrons and electrons.
The nucleus contains protons (+) and neutrons (0).
The electrons (-) are arranged in energy levels (orbitals) outside the nucleus.
In a neutral atom,
# protons = # electrons 8

9. The number of protons in the nucleus is the atomic number (Z).
The number of protons plus neutrons in the nucleus is the mass number (A).
Electrons of different atoms interact to form bonds. Therefore, the number of electrons an atom possesses determines its chemical behavior.
Size of atom (see fig 4.9)
Diameter of nucleus ~ cm
Diameter of overall atom ~ 10-8 cm
Diameter of an atom is about 100,000 times the diameter of the nucleus!
Atoms are mostly empty space!

10. Masses of parts of an atom
Electron = 1 (relative mass)
Proton = 1836
Neutron = 1839
-- most of the mass of an atom is located in the nucleus; the mass of the electrons is insignificant
In summary:
Atoms are very small, although the overall diameter of an atom is very large compared to the diameter of the atom’s nucleus.
Almost all the mass is located in the very tiny nucleus, making it very dense.

11. 4.7 Isotopes
4.7 Isotopes
All atoms of an element have the same number of protons
Atoms of an element with different numbers of neutrons are called isotopes
All isotopes of an element are chemically identical and behave the same in chemical reactions
Isotopes of an element have different masses
Isotopes are identified by their mass numbers
mass number = protons + neutrons 15

12. 4.8 Introduction to the Periodic Table
Isotope symbols
4.8 Introduction to the Periodic Table
See class handout for this section. 12 16

13. 4.9 Natural States of the Elements
4.9 Natural States of Elements
Elements are usually not found in pure form in nature (on earth).
Matter around us consists mostly of compounds and mixtures.
Most elements are reactive and form compounds with other elements.
Ex: Na + Cl  NaCl
Exceptions: the noble gases (Group 8A) and the noble metals (gold, silver, platinum).
Noble gases exist as single atoms.

14. Some gases exist as diatomic molecules at normal temperatures (around 25o C):H2
O2 and N2 are components of air
The halogens -- F2, Cl2, Br2, I2
However, hydrogen and the halogens are rarely present on earth in elemental form because they easily form compounds with other elements.
Ex: Na + Cl  NaCl
Mg + 2Br  MgBr2
2H2 + O2  2H2O

15. Only two elements are liquids at normal temperatures: bromine (Br2) and mercury (Hg) metal.
All other elements are solids at normal temperatures (around 25o C).
(cesium and gallium melt at about 30o C)

16. Allotropes
Forms of a solid nonmetallic element with different physical properties.
The different physical properties arise from the different arrangements of the atoms in the solid.
Allotropes of carbon include:
diamond (hard)
graphite (slippery)
buckminsterfullerene
large, soccer-ball-shaped molecules
See fig 4.18 16 21

17. 4.10 Ions
4.10 Ions
A neutral atom has zero net charge because the number of its protons (+) equals the number of its electrons (-).
Adding or removing electrons from an atom creates an ion – an atom with a net positive or negative charge.
When one or more electrons are lost by a neutral atom, an ion with a positive charge is formed, called an cation.
Examples:

18. A cation is named using the parent name of the atom:
Na+ is the sodium ion (or sodium cation)
Mg2+ is the magnesium ion (or magnesium cation)
When one or more electrons are gained by a neutral atom, an ion with a negative charge is formed, called an anion.
Examples:

19. An anion is named by taking the root name of the atom and adding the suffix –ide.
Cl- is the chloride ion
F- is the fluoride ion
O2- is the oxide ion
Ions are never formed by adding or removing protons to a nucleus.
Isolated atoms do not form ions on their own.

20. 4.11 Compounds That Contain Ions
Ions form when metallic elements react with nonmetallic elements. The metal atoms lose one or more electrons, which are in turn gained by the atoms of the nonmetal.

21. Ion Charges and the Periodic Table Group 1A metals form 1+ ions.
(See fig 4.19)
Group 1A metals form 1+ ions.
Group 2A metals form 2+ ions.
Group 3A metals form 3+ ions.
Transition metals form cations with various charges. Examples:
Fe2+ and Fe3+
Cu+ and Cu2+
Group 4A metals form cations with various charges:
Pb2+ and Pb4+
Sn2+ and Sn4+
Group 5A nonmetals form 3- ions.
Group 6A nonmetals form 2- ions.
Group 7A nonmetals form 1- ions