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

Pure Mixtur Heterogeneo Homogeneo Elemen Alloy Solutio 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 . 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. : the number of protons (positively charged particle) in the nucleus of an element. The number of electrons in an atom

2. 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 : : atomic number ______

metal/non-metal______

______

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: . 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: - “Salt Formers”

Reacts well with metals to form compounds similar to salts. Most diverse Group. Contains all phases. Contains most reactive non-metal:

Group 18: Noble Gases

Seldom reacts to form compounds. formerly called the “Inert Gases”

Group B elements: Transition Metals typical metals such as , , and silver wide variety of characteristics

Metal are seperated from the Non-Metals by the staircase Metal vs Non-Metal Properties

Metals Non-Metals -solid (except Hg) -s, l, g -silver(except: copper, ) -all colors -ductile & malleable -no -conduct heat/electricity -no -reacts with acid to form -some gas

Metalloids - 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 / Series: Name says it all, atomic #’s 58 - 71 Trans / 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 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 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 . 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 the atomic number (# of protons)

-protons equal 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 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: 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: 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 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. and 2. and 3. potassium and 4. barium and nitrogen

Notice a pattern? The charges of each element become the subscript for the other element. ie. Ca2+ P3- Ca P 3 2 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: borate 3- Li+ BO3 = Li3BO3 (s) If more than one complex ion is required then you must bracket the ion before writing the subscript.

Example: (II) chlorate 2+ 3- Ni ClO = Ni(ClO3)2 (s) Write the formula for each of the following: 1. Sodium carbonate 2. Potassium silicate 3. Magnesium hydroxide 4. (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 (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: dioxide - one carbon and two oxygen. di - means two. We must memorize a set of prefixes for naming these molecules.

mono - 1 hexa - 6 di - 2 hepta - 7 tri - 3 octa - 8 tetra - 4 nona - 9 penta - 5 deca -10

Example: CO - carbon monoxide SO3 - trioxide N2O - dinitrogen monoxide

Try: PCl5 - 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) water - H2O (l) ammonia - NH3 (g) methane - CH4 (g) methanol - CH3OH (l) butane - C4H10 (g) glucose - C6H12O6 (s) octane - C8H18 (l) propane - C3H8 (g) ethanol - C2H5OH (l) sucrose - C12H22O11 (s) hydrogen peroxide - H2O2 (l) hydrogen sulfide-H2S (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 (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)

H 2CO3 (aq)-carbonic acid (fizz in pop)

H 3PO4 (aq) -phosphoric acid (fertilizers)

H 2SO4 (aq) -sulfuric acid (battery acid)

HNO 3 (aq) -nitric acid (acid rain)

C 6H5COOH(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) NH 3 (g) - ammonia (household cleaner) Ca(OH)2 (s) - calcium hydroxide (Tums) NaHCO3 (s) - sodium bicarbonate (baking soda)