Download Chemistry Nomenclature Notes

Science 10 Chemistry Notes
Matter and Classification
Purpose of classification: to gain a better understanding and appreciate similarities or differences
Matter: Has mass and occupies space.
We can divide matter into two categories as follows:
Matte
Mixtur
Pure
Heterogeneo
Elemen
Homogeneo
Alloy Solutio
Metal Nonmeta
Compou
Ioni
Molecul
PURE SUBSTANCES
Substances that are the same or consistent throughout. Can be a single element or a combination of
elements.
Elements:
Substance composed of only one kind of atom. 109 on the periodic table. Each has a unique international
symbol. Can be combined to make other pure substances.
Compound:
Combination of two or more elements in specific proportions. Once combined the compound acts as one,
with consistent chemical and physical properties.
Chemical vs. Physical Properties
Chemical - describes the reactivity of a substance.
Physical -no new substance formed. Has similar properties in each form. Usually a change in phase. ie.
Liquid, solid or gas.
Chemical Reactions
The process that occurus whena subtance or substances reacts to creat a different substance.
- involve the production of new substances
- involve the flow of energy (exothermic and endothermic)
- involves formation of a gas
- involves the formation of a solid in a liquid
Periodic Table
Elements on the periodic table are classified and arranged according to four basic patterns:
1. Atomic number: the number of protons (positively charged particle) in the nucleus of an element.
The number of electrons in an atom
2. Metals vs. non-metals: separated by the staircase line.
3. Groups (or families): vertical columns that have similar properties.
4. Periods: horizontal rows which indicate the number of electron shells an atom has.
Example : Calcium:
atomic number __________
metal/non-metal___________
group ______________
period________________
The periodic table can also list the physical state (phase) of the element at room temperature.
Regular Print - solid
Clear Print - gas
Bold- liquid- There are only two (Hg, Br)
Chemical Families
Group 1 Alkali metals
Elements are highly reactive.
Contain most reactive metal: Francium.
Silver colored
Very ductile
React with air or water
Group 2 Alkaline Earth Metals
Similar to alkalies but not as reactive in air.
Oxidize with air to form a protective coating
Group 17: Halogens - “Salt Formers”
Reacts well with metals to form compounds similar to salts.
Most diverse Group. Contains all phases.
Contains most reactive non-metal: Fluorine
Group 18: Noble Gases
Seldom reacts to form compounds.
formerly called the “Inert Gases”
Group B elements: Transition Metals
typical metals such as copper, iron, zinc and silver
wide variety of characteristics
Metal are seperated from the Non-Metals by the staircase
Metal vs Non-Metal Properties
Metals
-solid (except Hg)
-silver(except: copper, gold)
-ductile & malleable
-conduct heat/electricity
-reacts with acid to form hydrogen gas
Non-Metals
-s, l, g
-all colors
-no
-no
-some Metalloids
-Nonmetals along the staircase line that have some of the properties of metals, mainly they can conduct
electricity. (semiconductors)
Last Two Groups - found at the bottom of the periodic table
Rare Earths / Lanthanide Series: Name says it all, atomic #’s 58 - 71
Trans Uranium /Actinide Series: Made in nuclear reactors, #’s 90-103
Atomic Structure
Atom: The basic unit of matter. Smallest unit of matter that retains the properties of the element.
Atom has a specific structure of subatomic particles consisting of:
1.
Proton: symbol “p+”
-Positively charged particle located in the center of an atom (nucleus). Makes up a large portion of
the mass of an atom.
2. Neutron: symbol “n°”
-Neutral charged particle located in the center of the atom(nucleus). Also makes up a large portion
of the mass of the atom
***Atomic mass of any element is determined by the number of protons and neutrons.
3.
Electron: symbol “e-”
-Negatively charged particle surrounding an atom.
-Has very little mass. Moves about the nucleus in an electron cloud. Cloud consists of mostly space.
If the nucleus was the size of a ping pong ball the first electron would be about 0.5 km away!
Finding Numbers of Subatomic Particles
Protons = Atomic Number
*Proton number can never, never, never change.
Electrons = Atomic Number (neutral atom only!) *Electron number can change.
Neutrons = Mass number - Atomic number.
Isotopes: Atoms that have a different number of neutrons from another atom of the same element.
U 235 is a light isotope having 3 fewer neutrons than the most common form of uranium
U 239 is a heavy isotope having 1 more neutron than the most common form of uranium
The Development of an Atomic Model
As theories developed old ideas were not discarded, they were modified and expanded upon.
Dalton “Billiard Ball” Model
Atom is in the shape of a billiard ball and acts as a single, indestructible and indivisible particle. The larger
the atomic number the larger the atom or “billiard ball”
Observations that supported this theory:
Law of Conservation of mass: The masses of the reactants always equals the masses of the products.
Example: 2 g of hydrogen and 16 g of oxygen would react to produce 18 g of water.
J.J. Thomson
“Raisin Bun” Model
Atoms have negatively charged particles embedded in them like raisins in a bun.
Observations that supported theory:
-electricity passed through a gas in a vacuum tube produced a stream of negatively charged particles.
Rutherford “Nuclear” Model
An atom’s mass is concentrated in a very small,dense and positively charged nucleus. Electrons orbit the
nucleus at a distance.
Observations that supported theory:
Gold foil experiment- large positively charged particles should go right through the gold foil. Most did but
some came right back towards the particle emitter.
Bohr
“Solar system” Model
Electrons are located in specific orbits, each having a specific energy level, around the nucleus. It is the
electrons in the outermost orbit that react with neighboring atoms to form compounds.
Observations that supported theory:
electricity passed through a gaseous element emits only certain wavelengths of light.
Quantum Mechanical Model
“Electron Cloud Model”
Electrons are in a cloud moving very quickly around a nucleus forming an electron cloud.
Atoms and Ions
Ions: charged particles (atoms) that have lost or gained electrons. They lose or gain electrons in order to
have an electron structure similar to that of a Noble gas.
Reason: Noble gases are stable!!
Comparing atoms to ions.
Atoms
Ions
-neutral charge
-positive or negative charge
-# of electrons equal
-# of electrons different from
to atomic number
(# of protons)
-protons equal atomic number
the atomic number -protons equal atomic number
Bohr Diagrams
# of protons goes in the nucleus
# of electrons can be distributed as follows:
maximum of 2 e- in the first level
maximum of 8 e- in the 2nd level
maximum of 8 e- in the third level
Draw a Bohr diagram for sodium and for fluorine.
METALS:
-Tend to lose electrons.
-They become positively charged and are called cations.
-The size of the positive charge is determined by the number of electrons lost.
-The number of electrons lost is determined by the proximity to the nearest Noble gas.
-Named by using the full metals name and adding ion at the end.
Ex: Magnesium is a group two element.
An atom of magnesium has ____ electrons.
The nearest Noble gas is ________ and it has ______ electrons. An ion of magnesium must also have
_______ electrons because this is a more stable configuration. This results in magnesium having a net
charge of 2+. It is named ___________________
NON-METALS:
-Tend to gain electrons.
-They become negatively charged and are called anions.
-The size of the negative charge is determined by the number of electrons gained.
-The number of electrons gained is determined by the proximity of Noble gas.
-Named by dropping the ending and adding an ‘ide’ ending
Ex: Chlorine is a group 7 element.
An atom of chlorine has _______ electrons.
The nearest Noble gas is ________ and it has ______ electrons. An ion of chlorine must also have
_______ electrons because it is more stable. This results in chlorine having a net charge of 1-. The name is
___
IONIC COMPOUNDS
Metals lose electrons to form positively charged ions called cations.
Nonmetals gain electrons to form negatively charged ions called anions.
When metals react with nonmetals an exchange of electrons occurs resulting in two oppositely charged ions.
It is these charges that cause the bond to form because opposite charges attract. The result is a crystal lattice.
Properties of ionic compounds:
- solid at room temperature
- dissolve in water (soluble) to form solution
- solutions conduct electricity (electrolytes)
- solutions can be any color
- have higher melting & boiling points
Charges must balance because one element gives up electrons and the other one accepts these same
electrons. The formula is the ratio of one ion to another.
Example 1: Sodium atoms tend to lose an electron to form the cation, Na1+. Chlorine atoms tend to gain
electrons to form an anion, Cl1-. When these two elements are brought together under the proper conditions
a chemical reaction takes place in which the sodium atom gives its electron to the chlorine atom. These two
ions attract each other and form a new compound, NaCl (s).
Name the compound by using the full name of the metal followed by the name of the nonmetal with the
‘ide’ ending. The above compound,NaCl (s), is named sodium chloride.
Sometimes more than one ion is required to react with another ion.
Example 2: Calcium reacts with fluorine. First, determine the charge on the ions. Look up each element on
the Table. Ca2+ and F1-. Calcium tends to lose two electrons but fluorine can only accept one. It takes two
fluoride ions to react with the calcium ion so the resulting compound is CaF2 (s). This means two fluoride
ions and one calcium ion form a compound. Name : calcium fluoride.
Always use the simplest whole numbers when writing the formula for an ionic compound.
Example 3: Zinc reacts with oxygen. Notice that the charges are equal in size. Zn2+ , O2-. The formula
for the compound is ZnO (s) and it is named zinc oxide.
Try these:
1. potassium and bromine
2. barium and iodine
3. potassium and nitrogen
4. barium and nitrogen
Notice a pattern? The charges of each element become the subscript for the other element. ie.
Ca2+ P3Ca3P2
Multivalent metals
Look at the transition metals. Notice that some of the elements show more than one charge is possible.
Example: iron 3+ on the left and 2+ on the right. When it reacts with a nonmetal such as oxygen it can
form two compounds, Fe2O3 (s) or FeO (s).
We need some way of distinguishing between the two compounds.
Name the metal followed by the size of the charge on the metal followed by the name of the nonmetal.
Example:
Fe2O3 (s) is named iron (III) oxide,
FeO (s) is named __________
Ensure that the charge is included in the name or it is wrong!!!
Polyatomic Ions
Notice the table in the upper portion labeled
“Table of Polyatomic Ions”. These are groups of atoms (mostly nonmetals) that form stable ions. They stay
together in most chemical reactions and are treated in the same way that individual ions are treated when
making an ionic compound. All rules for making formulas are followed.
Notice that most of the polyatomic ions are negatively charged except for ammonium. Also notice that the
names for most of the complex ions end in ‘ate’. This is your clue that the compound consists of a
polyatomic ion.
Example: Lithium borate
Li+ BO33- = Li3BO3 (s)
If more than one complex ion is required then you must bracket the ion before writing the subscript.
Example: nickel (II) chlorate
Ni2+
ClO3- = Ni(ClO3)2 (s)
Write the formula for each of the following:
1. Sodium carbonate
2. Potassium silicate
3. Magnesium hydroxide
4. Manganese (IV) phosphate
Name the following:
1. FeSO3 (s)
2. Cu2CO3 (s)
3. Co(IO3)2 (s)
4. Cr2(SO4)3 (s)
The table of complex ions only lists some of the possible ions. Many can be formed by adding an oxygen or
a hydrogen or by taking away an oxygen.
Example:
chlorate - ClO3- (most common ion)
perchlorate - ClO4- (one more oxygen)
chlorite - ClO2- (one less oxygen)
hypochlorite -ClO1- (two less oxygen)
Try:
Na3BO3 (s)
Na3BO2 (s)
Ca3(PO2)2(s)
*Note- the prefix “bi” means hydrogen!
Molecular Compounds
When two or more nonmetals react to form a compound, the result is a molecule. These molecules DO
NOT depend upon ionic charges. They are both negatively charged as ions so they would repel each other.
These atoms combine by sharing valence (outside) electrons. This type of bond is called covalent or
molecular. Since we cannot tell the formula from any charges the molecular substances are named
differently. The name of the molecule tells us the formula!
Example: carbon dioxide - one carbon and two oxygen. di - means two. We must memorize a set of
prefixes for naming these molecules.
mono - 1
di - 2
tri - 3
tetra - 4
penta - 5
Example: Try:
CO - carbon monoxide
SO3 - sulfur trioxide
N2O - dinitrogen monoxide
PCl5 -
hexa - 6
hepta - 7
octa - 8
nona - 9
deca -10
SF6 N2O4 -
Some molecules have become known by common names. These molecular names must be memorized.
Some of these have more than two non-metals. The ones you will be expected to know are:
ozone - O3 (g)
ammonia - NH3 (g) methanol - CH3OH (l)
butane - C4H10 (g)
glucose - C6H12O6 (s)
octane - C8H18 (l) ethanol - C2H5OH (l)
sucrose - C12H22O11 (s)
hydrogen peroxide - H2O2 (l)
hydrogen sulfide-H2S (g)
water - H2O (l)
methane - CH4 (g)
propane - C3H8 (g)
Most elements exist in compounds and must be refined to obtain pure substances. Some exist alone
(monatomic) such the Noble gases. Their chemical formula is just the elemental symbol. Others exist as
diatomic molecules. All group VIIA elements plus oxygen, nitrogen and hydrogen are diatomic. Their
formulas are as follows:
F2(g), Cl2(g), Br2(g), I2(g), At2(g), N2 ( g ) , O2(g), H2(g)
You must write the formulas for these molecules this way when they are alone.
There are two polyatomic molecules. They are sulfur (S8 (s)) and phosphorus (P4(s)).
Acids and Bases
Acids - All have pH values of less than 7.
Properties:
1. soluble in water
2. solutions conduct electricity(electrolytes)
3. react with metals to produce hydrogen gas
4. taste sour
5. neutralize bases
6. turn blue litmus red
Naming Acids
Steps:
1. First you must identify the compound as being an acid:
A) Has aqueous as the phase (aq)
B) There is Hydrogen in the formula (usually at the start of the formula)
2. Name it as an ionic compound!
3. Then use the chart below to convert it from an ionic name to an acid name
*The ending of the ionic name determines the acid name
hydrogen _____ide becomes hydro_____ic acid
hydrogen _____ate becomes _________ic acid
hydrogen _____ite becomes ________ous acid
To use the table you must fill in the blanks with a root word.
Example:
1.
hydrogen chloride -->hydrochloric acid
2.
hydrogen carbonate -->carbonic acid
3.
hydrogen sulfite --> sulfurous acid
Note: elements sulfur and phosphorus usually use the entire name when naming an acid compound.
Reason- it just sounds better.
Common acids:
HCl (aq) -hydrochloric acid (stomach acid)
H2CO3 (aq)-carbonic acid
(fizz in pop)
H3PO4 (aq) -phosphoric acid (fertilizers)
H2SO4 (aq) -sulfuric acid (battery acid)
HNO3 (aq) -nitric acid (acid rain)
C6H5COOH(aq) -benzoic acid (preservative)
CH3COOH (aq) -acetic acid (main part of vinegar)
Bases:
Compounds that have a pH of more than 7 and that usually dissolve in water to form hydroxide ions (OH-).
Properties:
1. soluble in water
2. solutions conduct electricity (electrolytes)
3. feel slippery to the touch
4. neutralize acids
5. cause red litmus to turn blue
Some common bases are:
NaOH (s) - sodium hydroxide (soap)
Mg(OH)2 (s)- magnesium hydroxide (Milk of Magnesia, Rolaids)
NH3 (g) - ammonia (household cleaner)
Ca(OH)2 (s) - calcium hydroxide (Tums)
NaHCO3 (s) - sodium bicarbonate (baking soda)