Domain 2: Chemistry- Chemical Reactions and Properties of ...
Transcript of Domain 2: Chemistry- Chemical Reactions and Properties of ...
Matter• Matter is anything that: a) has
mass, and b) takes up space
• Mass = a measure of the amount
of “stuff” (or material) the object
contains (don’t confuse this with
weight, a measure of gravity)
• Volume = a measure of the space
occupied by the object
Properties are…• Words that describe matter (adjectives)
• Physical Properties- a property that can be observed and measured without changing the material’s composition.
• Examples- color, hardness, m.p., b.p.
• Chemical Properties- a property that can only be observed by changing the composition of the material.
• Examples- ability to burn, decompose, ferment, react with, etc.
States of matter1) Solid- matter that can not flow (definite
shape) and has definite volume.
2) Liquid- definite volume but takes the shape of its container (flows).
3) Gas- a substance without definite volume or shape and can flow.
– Vapor- a substance that is currently a gas,
but normally is a liquid or solid at room
temperature. (Which is correct: water gas,
or water vapor?)
States of MatterStates of Matter
Solid
Liquid
Gas
Definite Volume?
YES
YES
NO
Definite Shape?
YES
NO
NO
Result of a
TemperatureIncrease?
Small
Expans.
Small
Expans.
Large
Expans.
Will it Compress?
NO
NO
YES
Physical vs. Chemical Change• Physical change will change the
appearance, without changing the
composition of the material.
– Boil, melt, cut, bend, split, crack
– Is boiled water still water?
• Can be reversible, or irreversible
• Chemical change - a change where a
new form of matter is formed.
– Rust, burn, decompose, ferment
• Mixtures are a physical blend of at least two substances; have variable composition. They can be either:
1) Heterogeneous – the mixture is not uniform in composition
• Chocolate chip cookie, gravel, soil.
2) Homogeneous - same composition throughout; called “solutions”
• Kool-aid, air, salt water
• Every part keeps it’s own properties.
Separating Mixtures
• Some can be separated easily by
physical means: rocks and marbles, iron filings and sulfur (use magnet)
• Differences in physical properties
can be used to separate mixtures.
• Filtration - separates a solid from the liquid in a heterogeneous mixture (by size) – Figure 2.7, page 46
Substances• Elements- simplest kind of matter
– cannot be broken down any simpler
– all one kind of atom.
• Compounds are substances that can be broken down only by chemical methods– when broken down, the pieces have
completely different properties than the original compound.
– made of two or more atoms, chemically combined (not just a physical blend!)
Compound or Mixture?
Compound Mixture
Made of one kindof material
Made of more thanone kind of material
Made by a chemical change
Made by a physical change
Definitecomposition
Variablecomposition
Elements vs. Compounds
• Compounds can be broken down
into simpler substances by chemical means, but elements
cannot.
• A “chemical change” is a change that produces matter with a different composition than the original matter.
Naming cations
• Two methods can clarify when
more than one charge is possible:
1) Stock system – uses roman
numerals in parenthesis to indicate the numerical value
Naming cations• We will use the Stock system.
• Cation - if the charge is always the same (like in the Group A metals) just write the name of the metal.
• Transition metals can have more than one type of charge.
• Indicate their charge with roman numerals in parenthesis after the name of the metal
Naming Anions
• Anions are always the same charge
• Change the monatomic element ending to – ide
• F1- a Fluorine atom becomes a Fluoride ion.
Polyatomic ions are…• Groups of atoms that stay together and
have an overall charge, and one name.
• Usually end in –ate or -ite
• Acetate: C2H3O21-
• Nitrate: NO31-
• Nitrite: NO21-
• Permanganate: MnO41-
• Hydroxide: OH1- and Cyanide: CN1-?
Writing Ionic Compound Formulas
Example: Barium nitrate
1. Write the formulas for the cation and anion, including CHARGES!
BaBa2+2+ NONO33--
2. Check to see if charges are balanced.
3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Use the criss-cross
method to balance subscripts.
Not balanced!
( )( )22
Now balanced.
= Ba(NO3)2
Writing Ionic Compound Formulas
Example: Iron (III) chloride
1. Write the formulas for the cation and anion, including CHARGES!
FeFe3+3+ ClCl--
2. Check to see if charges are balanced.
3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Use the criss-cross
method to balance the subscripts.
Not balanced!
33Now balanced.
= FeCl3
Writing Ionic Compound Formulas
Example: Magnesium carbonate
1. Write the formulas for the cation and anion, including CHARGES!
MgMg2+2+ COCO3322--
2. Check to see if charges are balanced.
They are balanced!
= MgCO3
Naming Ionic CompoundsNaming Ionic Compounds• 1. Name the cation first, then anion
• 2. Monatomic cation = name of the
element
Ca2+ = calcium ion
• 3. Monatomic anion = root + -ide
Cl− = chloride
CaCl2 = calcium chloride
Naming Ionic CompoundsNaming Ionic Compounds
• some metals can form more than one
charge (usually the transition metals)
• use a Roman numeral in their name:
PbCl2 – use the anion to find the charge on the cation (chloride is always 1-)
Pb2+ is the lead (II) cation
PbCl2 = lead (II) chloride
(Metals with multiple oxidation states)(Metals with multiple oxidation states)
Things to look for
• If cations have (_), the number
in parenthesis is their charge.
• If anions end in -ide they are
probably off the periodic table
(Monoatomic)
• If anion ends in -ate or –ite,
then it is polyatomic
Molecular compounds are…
• made of just nonmetals
• smallest piece is a molecule
• can’t be held together
because of opposite charges.
• can’t use charges to figure out how many of each atom
Molecular compounds are easier!• Ionic compounds use charges to
determine how many of each.
–Have to figure out charges.
–May need to criss-cross numbers.
• Molecular compounds: the name
tells you the number of atoms.
• Uses prefixes to tell you the exact
number of each element present!
Prefixes (Table 9.4, p.269)• 1 = mono-
• 2 = di-
• 3 = tri-
• 4 = tetra-
• 5 = penta-
• 6 = hexa-
• 7 = hepta-
• 8 = octa-
Prefixes
• 9 = nona-
• 10 = deca-
• To write the name, write two words:
• One exception is we don’t write mono- if there is only one of the first element.
• Normally do not have double vowels when writing names (oa oo)
Prefix name Prefix name -ide
Write formulas for these:• diphosphorus pentoxide
• tetraiodine nonoxide
• Iron (III) oxide
• nitrogen trioxide
• Calcium sulfide
• Potassium iodide
• aluminum chloride
Chemical Reactions are…• When one or more substances are
changed into new substances.
• Reactants- the stuff you start with
• Products- what you make
• The products will have NEW
PROPERTIES different from the
reactants you started with
• Arrow points from the reactants to the
new products
Recognizing Chemical Changes
1) Energy is absorbed or released
(temperature changes hotter or colder)
2) Color changes
3) Gas production (bubbling, fizzing, or odor
change; smoke)
4) formation of a precipitate - a solid that
separates from solution (won’t dissolve)
5) Irreversibility - not easily reversed
But, there are examples of these that are not
chemical – boiling water bubbles, etc.
Conservation of Mass • During any chemical reaction, the
mass of the products is always equal
to the mass of the reactants.
• All the mass can be accounted for:
– Burning of wood results in products that appear to have less mass as ashes; where is the rest?
• Law of conservation of mass
4. Balanced Chemical Equations
• Atoms can’t be created or destroyed in an ordinary reaction:
–All the atoms we start with we must end up with
• A balanced equation has the
same number of each element on
both sides of the equation.
Rules for balancing:1) Assemble the correct formulas for all the
reactants and products, use + and →
2) Count the number of atoms of each type appearing on both sides
3) Balance the elements one at a time by adding coefficients where needed (the numbers in front) - save balancing the H and O until LAST!
(I prefer to save O until the very last)
4) Check to make sure it is balanced.
• Never change a subscript to balance an equation.
– If you change the formula you are describing a different reaction.
– H2O is a different compound than H2O2
• Never put a coefficient in the middle of a formula
2NaCl is okay, but Na2Cl is not.
Types of Chemical Reactions
• Combination A+B � AB
• Decoposition AB� A + B
• Single Replacement AB + C � AC + B
or A + BC � AC + B
• Double Replacement AB + CD � AD + CB
• Combustion CH(O) + O2 � CO2 + H20
• If O2 limited CH(O) + O2 � CO + H20
• If O2 limited CH(O) + O2 � C + H20
#1 - Combination Reactions• Combine = put together
• 2 substances combine to make one compound.
• Ca +O2 → → → → CaO
• SO3 + H2O →→→→ H2SO4
• We can predict the products if the reactants are two elements.
• Mg + N2 → ____→ ____→ ____→ ____
#2 - Decomposition Reactions• decompose = fall apart
• one reactant breaks apart into two or more elements or compounds.
• NaCl Na + Cl2• CaCO3 CaO + CO2
• Note that energy (heat, sunlight,
electricity, etc.) is usually required
electricity →→→→
∆∆∆∆ →→→→
#3 - Single Replacement
• One element replaces another
• Reactants must be an element and a compound.
• Products will be a different element and a different compound.
• Na + KCl →→→→ K + NaCl
• F2 + LiCl →→→→ LiF + Cl2
#4 - Double Replacement• Two things replace each other.
– Reactants must be two ionic compounds.
– Usually in aqueous solution
• NaOH + FeCl3 →
– The positive ions change place.
• NaOH + FeCl3 → Fe+3 OH- + Na+1 Cl-1
• NaOH + FeCl3 → Fe(OH)3 + NaCl
#5 - Combustion• Means “add oxygen”
• Normally, a compound composed of only C, H, (and maybe O) is reacted with oxygen – usually called “burning”
• If the combustion is complete, the products will be CO2 and H2O.
• If the combustion is incomplete, the products will be CO (or possibly just C) and H2O.
Solvents and Solutes• Solution - a homogenous mixture, that
is mixed molecule by molecule.
• Solvent - the dissolving medium
• Solute -the dissolved particles
• Aqueous solution- a solution with water
as the solvent.
• Particle size less than 1 nm; cannot be
separated by filtration – Fig. 15.6, p.450
1. SoluteA solute is the dissolved substance in a
solution.
A solvent is the dissolving medium in a
solution.
2. Solvent
Salt in salt water Sugar in soda drinks
Carbon dioxide in soda drinks
Water in salt water Water in soda
Parts of a Solution:
Dissolution of Solid SoluteWhat are the driving forces which cause solutes to dissolve to form solutions?
1. CovalentCovalent solutes dissolve by H-bonding to water
2. Ionic solutes dissolve by dissociation into
their ions.
These ions have
been surrounded
by water, and are
now dissolved!
• Solids will dissolve if the attractive force of the water molecules is stronger than the attractive force of the crystal.
• If not, the solids are insoluble.
• Water doesn’t dissolve nonpolarmolecules (like oil) because the water molecules can’t hold onto them.
• The water molecules hold onto other water molecules, and separate from the nonpolar molecules.
• Nonpolars? No repulsion between them
Electrolytes and Nonelectrolytes
• Electrolytes- compounds that conduct
an electric current in aqueous solution,
or in the molten state
– all ionic compounds are electrolytes because they dissociate into ions (they are also called “salts”)
• barium sulfate- will conduct when
molten, but is insoluble in water!
Electrolytes and Nonelectrolytes• Do not conduct? = Nonelectrolytes.
– Most are molecular materials, because they do not have ions
• Not all electrolytes conduct to the
same degree
– there are weak electrolytes, and strong electrolytes
– depends on: degree of ionization
Solution formation
• The “nature” (polarity or composition) of the solute and the solvent will determine…
1. Whether a substance will dissolve
2. How much will dissolve
• Factors determining rate of solution...
1. stirring (agitation)
2. surface area the dissolving particles
3. temperature
Making solutions• In order to dissolve, the solvent
molecules must come in contact
with the solute.
1. Stirring moves fresh solvent into
contact with the solute.
2. Smaller pieces increase the amount
of surface area of the solute.
- think of how fast a breath mint
dissolves when you chew it
Temperature and Solutions
3. Higher temperature makes the
molecules of the solvent move around faster and contact the solute harder and more often.
– Speeds up dissolving.
• Higher Temperature Also Usually
increases the amount that will dissolve (an exception is gases)
Properties of acids
• Taste sour (don’t try this at home).
• Conduct electricity.
–Can be strong or weak electrolytes in
aqueous solution
• React with metals to form H2 gas.
• Change the color of indicators (blue litmus to red).
• React with bases (hydroxides) to form
water and a salt.
Acids Neutralize Bases
HCl + NaOH → NaCl + H2O
-Neutralization reactions ALWAYS produce a salt and water.
-Of course, it takes the right proportion of acid and base to produce a neutral salt
Properties of bases• React with acids to form water
and a salt.
• Taste bitter.
• Feel slippery (don’t try this either).
• Can be strong or weak electrolytes in aqueous solution
• Change the color of indicators (red litmus turns blue).
Bases Affect Indicators
Red litmus paper
turns blue in contact
with a base.Phenolphthalein
turns purple in a
base.