1. Chapter 1 Notes Atoms and Bonding

2. Section 1: Elements and Atoms
A. Matter is defined as anything that has mass and takes up space.
- Ancient Greeks believed that all matter was made up of four elements; earth, air, fire and water.

3. B .Elements are the simplest pure substances, and they cannot be broken down into any other substance. (the periodic table of elements)
- Often called the building blocks of matter because all matter is composed of one element or a combination of two or more elements.
C. Compound – a pure substance made of two or more elements that are chemically combined.
Ex. Sodium Chloride (table salt) = NaCl

4. D. Mixture – two or more substances that are in the same place but are not chemically combined.
Ex. Air, soil, wood, orange juice

5. E. Atoms
1. Democritus (430 B.C.) Proposed the idea of an atom.
2. Matter is formed of pieces that are “uncuttable”
3. Modern definition:
a. An atom is the smallest particle of an element.

6. 2. Atomic Theory and Models
A scientific theory- a well-tested idea that explains and connects a wide range of observations
Daltons atomic theory-
All elements are composed of atoms.
All atoms of the same element are exactly alike and have the same mass. Atoms of different elements are different and have different masses.
An atom of one element cannot be changed into an atom of a different element. Atoms cannot be created or destroyed in any chemical change, only rearranged.
Every compound is composed of atoms of different elements, combined in a specific ratio.

7. Found atoms contained negatively charged particles called electrons.
Thomson-
Found atoms contained negatively charged particles called electrons.
Aka, “muffin model”

8. Rutherford- gold foil experiment
Inferred that an atoms positive charge must be clustered in a tiny region in its center called the nucleus
The nucleus is made
of one or more positively
charged particles called
protons

9. Bohr’s Model-electrons orbit the nucleus in certain orbits similar to the way planets orbit the sun.

10. Electron Cloud Model-electrons can be anywhere in a cloudlike region around the nucleus. The “cloud” symbolizes where an electron is likely to be found. An electron’s movement is related to its energy level- specific amount of energy it has. Electrons with more energy are farther from the nucleus.
Modern Atomic Model-James Chadwick discovered the neutron, a neutral particle in the nucleus with nearly the same mass as a proton.

11. Section 2: Atoms, Bonding, and the Periodic Table

12. 1. Valence electrons and Bonding
a. valence electrons- electrons that are in the highest energy level of an atom and are held loosely by the nucleus
-valence electrons determine the properties of the element and how the atom can bond with other elements

13. b. electron dot diagrams-symbol of an element surrounded by the number of dots that equals the number of valence electrons the element has.

14. c. Chemical bond- the force of attraction that holds two atoms together as a result of the rearrangement of electrons between them
-atoms are less likely to react with they have 8 valence electrons.
-atoms will gain or lose electrons to become more stable.
-when atoms bond, electrons can be transferred from one atom to another or are shared.
-the bonding of atoms causes a chemical reaction and makes a new substance.

15. 2. The Periodic Table
A. First put together in the 1800’s by a Russian Chemist –Dimitri Mendeleev
B. Modern periodic table is used worldwide
1. Elements are organized into groups with the same number of valence electrons in their atoms.
2. Important information contained within each block

17. a. atomic number = number of protons
1) protons are positively charged subatomic particles
2) nucleus is found in the center of the atom. It is made up of protons and neutrons.
3) neutron – is a neutral particle – nearly the same mass as a proton.
4) electron – typically equals the number of protons in the Atom, except if the atom is an ion, it is a negative particle that orbits around the nucleus in different energy levels.
5). atomic mass = the number of protons and neutrons
in the nucleus of an atom.

18. Calculations involving Subatomic Particles:
atomic number = # of protons
mass number = # of protons + # of neutrons
neutral atom: # of protons = # of electrons
charged ion: # of electrons = # of protons - +charge
# of electrons = # of protons + - charge

19. 3. Organization of the periodic table (see periodic table handout)
Groups (family)– a column (vertical) of elements that have the same number of valence electrons.
Period – elements in the same horizontal row of the periodic table.
Elements are classified into three major categories that include metals, nonmetals, and metalloids.

20. 4. Families of the periodic table
a. alkali metals-very reactive; Li, Na, K, Rb, Cs, Fr
b. alkaline earth- reactive; Be, Mg, Ca, Sr, Ba, Ra
c. halogens-“salt formers”; F, Cl, Br, I
d. noble gases- all stable, few react; He, NE, Ar, Kr, Xe, Rn

21. Section 3 Ionic Bonds

22. 1. Ions and ionic bonds a. ion- an atom or group of atoms with a charge. -cation- an atom that loses an electrons and has a positive charge. -anion- an atom that gains an electrons and has a negative charge. b. polyatomic ion- ions that are made of more than one atom and can be positive or negative. c. ionic bonds-attraction between oppositely charged ions. d. an ionic compound is one that consists of positive and negative ions. e. cations are usually metals and anions are usually nonmetals

23. 2. Chemical formulas and names
Formula writing
Ex: sodium chloride NaCl
Magnesium chloride  MgCl2

24. Formula units have a net zero charge Determining Charge
Formula Unit-Simplest ratio of the ions represented in an ionic compound.
Formula units have a net zero charge
Determining Charge
monoatomic ions –a one atom ion
Charge based on the location of the element on the periodic table
3. oxidation number –charge of the monatomic ion 1A 1+ 6A 2- 2A 2+ 7A 1- 5A 3- 8A

25. Oxidation numbers

26. Determining the Formula :
Criss-cross method
Cations listed first and anions listed second
Find the symbols and charges in the upper right corner of the symbol
3. Cross the charges, drop the (+) or (-) sign
4. Reduce if necessary and never write the number 1

27. Examples:
potassium oxide ______________________
aluminum sulfide ______________________
calcium fluoride _______________________
zinc fluoride ________________________
aluminum iodide _______________________

28. Polyatomic ions (see handout for list)
1. Write the symbols and charges
2. Cross the charges and drop the sign
3. If the oxidation number going to the polyatomic is greater than 1, add parenthesis.
4. Reduce the charges if possible, but do not change the polyatomic

29. Naming Ionic Compounds (continued)
II. Polyatomic ions each have specific names which must be memorized so they can be recognized on sight.
(At this point, if you are asked to name any compound that contains more than two elements, it will contain at least one polyatomic ion.)
A few of the more common polyatomic ions
Formula
Name
C2H3O21-
acetate
CO32-
carbonate
HCO31-
bicarbonate
NH41+
ammonium
Formula
Name
NO31-
nitrate
OH1-
hydroxide
PO43-
phosphate
SO42-
sulfate

30. Examples: sodium nitrate ______________________
potassium acetate ___________________
aluminum sulfate_____________________
lithium chlorate ______________________

31. Naming ions and ionic compounds
1. oxyanions –polyatomic ion composed of an element, usually a metal and oxygen
The ion with more oxygen atoms is named using the root of the nonmetal plus the suffix -ate.
The ion with fewer oxygen atoms is named using the root of the nonmetal plus the suffix -ite

32. Rules for naming. 1)Name the cation first.
2)Monoatomic cations use the element name
3)Monoatomic anions take their name from the root of the element name and add –ide
4)Some metals can have multiple oxidation numbers – distinguish them by using Roman numerals
Ex Fe and Fe+3
iron (II) iron (III)
5)Group I metals always +1 and Group II metals always +2
6)If the compound contains a polyatomic ion, simply name the ion.

34. Naming Ionic Compounds (continued)
If a Roman numeral is required, the charge on the metal ion must be determined from the charge on the negative ion.
Helpful Rules to Remember
A metal ion is always positive.
The Roman numeral indicates the charge, not the subscript.
The positive and negative charges must cancel (total charge must = 0).
Nonmetals are always negative & can never form more than one monatomic ion.
Examples
Formula
Reasoning
Name
FeCl2
Cl has a 1- charge, and there are 2 of them for a total of 2-, so the Fe must be 2+
iron (II) chloride
Fe2O3
O has a 2- charge, and there are 3 of them for a total of 6-, so the Fe must have a total charge of 6+ split equally between the two iron atoms, so each must have a 3+ charge
iron (III) oxide
PbS2
S has a 2- charge, and there are 2 of them for a total of 4-, so the Pb must be 4+
lead (IV) sulfide
Cu3N
N has a 3- charge, so the Cu must have a total charge of 3+ split equally between the 3 copper atoms, so each must have a 1+ charge
copper (I) nitride

35. Additional Examples: C2H3O21- CO32- HCO31- NH41+ Na2SO4 Fe(NO3)2 AlCl3
PbI4
(NH4)3PO4
Mg3N2
C2H3O21-
acetate
CO32-
carbonate
HCO31-
bicarbonate
NH41+
ammonium
NO31-
nitrate
OH1-
hydroxide
PO43-
phosphate
SO42-
sulfate
* Groups I & II, Al, Zn, Cd, and Ag need no Roman numeral.

36. AgC2H3O2 ________________ NaOH ___________________
NaClO3 ___________________
NaC2H3O2 _________________
BaSO4 ___________________
ZnCO3 ___________________
C2H3O21-
acetate
CO32-
carbonate
HCO31-
bicarbonate
NH41+
ammonium
NO31-
nitrate
OH1-
hydroxide
PO43-
phosphate
SO42-
sulfate
* Groups I & II, Al, Zn, Cd, and Ag need no Roman numeral.

37. Naming Ionic Compounds: Examples
C2H3O21-
acetate
CO32-
carbonate
HCO31-
bicarbonate
NH41+
ammonium
NO31-
nitrate
OH1-
hydroxide
PO43-
phosphate
SO42-
sulfate
* Groups I & II, Al, Zn, Cd, and Ag need no Roman numeral.
Na2SO4
sodium sulfate
Fe(NO3)2
iron (II) nitrate
AlCl3
aluminum chloride
PbI4
lead (IV) iodide
(NH4)3PO4
ammonium phosphate
Mg3N2
magnesium nitride
AgC2H3O2
silver acetate

38. Properties of Ionic Compounds
a. hard, brittle crystals b. have high melting points c. solids at room temperature d. conduct electricity when dissolved in water

39. A covalent bond is one that results from the sharing of valence electrons
A molecule is formed when two or more atoms bond covalently
The difference between an ionic compound and a molecule is that an ionic compounds donate electrons and molecules share valence electrons
9.1 Covalent Bonds

40. Section 4 Covalent BOnds

41. How Covalent bonds form
How Covalent bonds form
When two atoms share electrons it is called a covalent bond
Covalent bonds form between nonmetals
A molecule is a neutral group of atoms formed by covalent bonds
Molecules can have single, double or triple bonds

42. 2. Naming Molecular Compounds
1. The first element in the formula is always named first, using the entire element name.
2. The second element in the formula is named using the root of the element and adding the – ide
3. Prefixes are used to indicate the number of atoms of each type that are present (Exception: never use a prefix if there is only one of the first element)

43. Naming Covalent Compounds
Prefixes
Subscript
Prefix 1
mono- 2
di- 3
tri- 4
tetra- 5
penta-
Subscript
Prefix 6
hexa- 7
hepta- 8
octa- 9
nona- 10
deca-
Note: When a prefix ending in ‘o’ or ‘a’ is added to ‘oxide’, the final vowel in the prefix is dropped.

44. 3. Writing Molecular Formulas
A) Write the symbols for the elements
B) Place the number of the prefix after the element symbol
C) Do not cross charges or reduce

45. Examples: N2S4 NI3 XeF6 CCl4 P2O5 SO3 1 mono 2 di 3 tri 4 tetra 5
penta 6
hexa 7
hepta 8
octa 9
nona 10
deca
* Second element in ‘ide’ from
* Drop –a & -o before ‘oxide’
N2S4
NI3
XeF6
CCl4
P2O5
SO3

46. Naming Binary Covalent Compounds: Examples1
mono 2 di 3
tri 4
tetra 5
penta 6
hexa 7
hepta 8
octa 9
nona 10
deca
* Second element in ‘ide’ from
* Drop –a & -o before ‘oxide’
N2S4
dinitrogen tetrasulfide
NI3
nitrogen triiodide
XeF6
xenon hexafluoride
CCl4
carbon tetrachloride
P2O5
diphosphorus pentoxide
SO3
sulfur trioxide

47. Examples: 1 mono 2 di 3 tri 4 tetra 5 penta 6 hexa 7 heptaa 8 octa 9
nona 10
deca
* Second element in ‘ide’ from
* Drop –a & -o before ‘oxide’
xenon trioxide _____________________________
bromine trichloride _____________________________
nitrogen trifluoride _____________________________
tetrasulfur tetranitride _____________________________
ditellurium dichloride _____________________________
dinitrogen tetraoxide _____________________________
phosphorus pentachloride____________________________

48. xenon trioxide XeO3
bromine trichloride BrCl3
nitrogen trifluoride NF3
tetrasulfur tetranitride S4N4
ditellurium dichloride Te2Cl2
dinitrogen tetraoxide N2O4
phosphorus pentachloride PCl5

49. b. have low melting and boiling points c. most are gases and liquids
3. Properties of Molecular compounds
a. molecular compound- compound of molecules that are covalently bonded
b. have low melting and boiling points
c. most are gases and liquids
d. do not conduct electricity