1. Atoms and Molecules, Inside and Out
Chemistry 1405
Chapters 5

2. Atoms
An atom is the smallest unit of matter that has the properties of an element.
The atom contains:
Electrons – found outside the nucleus; negatively charged.
Protons – found in the nucleus; positive charge equal in magnitude to the electron’s negative charge.
Neutrons – found in the nucleus; no charge; virtually same mass as a proton.
Mass of atom is in the nucleus; nucleus is small compared to electron cloud.

3. Atomic Theory of Matter
All matter is made of atoms. 118 different types of atoms known.
The atoms of a given element are identical; the atoms of different elements are different in some fundamental way or ways.
Relative number and arrangement of different types of atoms in a pure substance determine its identity.

4. Atomic Theory of Matter
Chemical change is a union, separation or rearrangement of atoms to give new substances.
Only whole atoms can participate in or result from any chemical change.
For our purposes atoms are considered indestructible during such changes.

5. Atoms Atoms are neutral
Atoms will have the same number of protons and electrons. A balanced charge means they are neutral.
Ions are not neutral.
Atoms make up elements and compounds; Elements are pure substances with all the same atomic number.
Compounds chemical are combinations of atoms.

6. Atomic Number(Z) The number of protons in the nucleus of an atom.
Atoms are defined by their number of protons and neutrons they contain.
In a neutral atom the number of protons is equal to the number of electrons, so the atomic number also indicate the number the number of electrons present in an atom.
For example, the atomic number of fluorine is 9, This means that each fluorine atom has 9 protons and 9 electrons.

7. Mass number (A)
Sum of the protons and neutrons in the nucleus of an atom.
The number of neutrons in an atom is equal to the difference between the mass number and the atomic number ( A-Z)
Example: Aluminum has 13 protons and 14 neutrons so the mass number is 27.

8. Nuclear Notation
Nuclear Notation: used to indicate the atomic number and mass number of an atom or isotope
Nuclear Notation is formed by writing an elemental symbol preceded by a subscript indicating its atomic number (Z) and a superscript indicating its mass number (A).
Nuclide:
An atom with a specific atomic number and a specific atomic mass number.

9. Isotopes All isotopes of an element are chemically identical
undergo the exact same chemical reactions
All isotopes of an element have the same number of protons
Isotopes of an element have different masses
Isotopes of an element have different numbers of neutrons
Isotopes are identified by their mass numbers, which is the sum of all the protons and neutrons in the nucleus
the percentage of an element that is one isotope is called the isotope’s natural abundance

10. Isotopes of potassium
The three naturally occurring potassium isotope

11. Isotopes of carbon 12, 13, and 14

12. Isotopes Percent Abundance:
Percent of atoms in a natural sample of a pure element that are a particular isotope of the element. (See table on p. 157 of 11th edition)
Distribution of isotopes in any particular sample is generally constant.
Is used to determine an atomic mass unit (amu)
23 elements have only one naturally occurring form. (Be, F, Na, Al, P)
Most naturally occurring elements are mixtures of isotopes.

14. Isobars Isobars: Atoms that have the same mass number.
They will have different atomic numbers.
58Fe 58Ni 26 28

15. Using the information below, indicate if the pairs of atoms are isotopes, isobars or neither.
Protons
Neutrons
Electrons
Atom A 9 10
Atom B
Atom C
Atom D
Atom A and Atom B
Atom C and Atom D
Atom A and Atom D
Atom A and Atom C
Atom C and Atom B
Atom B and Atom D

16. Mass number – Atomic mass
Mass number = sum of protons and neutrons
Atomic mass = weighted average mass that accounts for:
The number of isotopes for the element.
percent abundance of each isotope.
the relative mass of each isotope using carbon as the reference point.
mass variations of the isotopes of a particular element.

17. Molecules
Molecules are made from 2 or more atoms covalently bound together which function as a single entity.
These are called molecular compounds.
The relative number and arrangement of different types of atoms contained in a pure substance determine its identity.
The molecule is the limit of physical subdivision.
The atom is the limit of chemical subdivision.

18. Ions
An ion is an atom or a group of atoms that have a positive or negative charge.
An ion forms when an atom or group of atoms loses or gains electron.
The loss of one or more electrons from a neutral atom result in a cation, an ion with a positive charge.
On the other hand, an anion is an ion whose net charge is negative due to an increase in the number of electrons.

19. Example of Ions
Sodium Atom (Na)

20. Molecules are classified into two type.
A homoatomic molecule is a molecule in which all atoms present are the same kind
A heteroatomic molecule is a molecule in which two or more different kinds of atoms are present.

21. Molecules are classified into two type.
A homoatomic molecule is a molecule in which all atoms present are the same kind
A heteroatomic molecule is a molecule in which two or more different kinds of atoms are present.

22. Examples of homo and hetero atomic molecule

23. Molecules
Chemical formulas are used to specify the types different atoms and how many of each atom make up covalent and ionic compounds.
Example Helium and Neon have the chemical formula He and Ne
A molecule is the smallest unit of a compound (covalent compound).
A formula unit is the smallest unit of a formula unit (ionic compound).
Remember, molecular formulas and formula units represent covalent and ionic molecules that make up compounds.

24. Chemical Formulas
A chemical formula is a combination of symbols that represents the composition of a compound.
Note: if the chemical symbol uses two letters, only the first letter is capitalized.
CoCl2 vs COCl2 Cobalt (II) chloride vs phosgene gas
They use subscripts to tell how many of each type of atom are present in the formula.
If only one atom is in the formula, no subscript is used and one is assumed.
If more than one of a group of atoms is present, parenthesis are used.