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SCIENCE10:CHEMISTRY1
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Atomic Theory and Bonding
The Nucleus
- The particles that make up an atom are called subatomic particles
- The three subatomic particles are protons, neutrons and electrons.
- Protons, which have a +1 (positive) electric charge, and neutrons which do not have an electric charge, are
located in the center of the atom or nucleus of each atom.
- The electric charge on any nucleus is always positive since protons are positively charged and neutrons are
neutral
- The charge on the nucleus is called the nuclear charge
- The nuclear charge is therefore equal to the number of protons (which is also equal to the atomic number)
- Electrons, which have a -1 (negative) electric charge surround the nucleus in regular patterns called orbitals or
energy levels.
Refer to the Table below in order to answer the following questions.
Name Symbol Electric Charge
Location in the Atom
Relative Mass
Proton p +1 Nucleus 1836 Neutron n 0 Nucleus 1837 Electron e -1 Surrounding the
nucleus 1
Ø Which subatomic particle or particles account for the majority of the volume of an atom? electrons Ø Which subatomic particle or particles account for the majority of the mass of an atom? Protons AND
neutrons
Organization of the Periodic Table
• Each element is listed according to its atomic number
• Each row is called a period ___________________
• Each colomn is called a group or family . Elements within a family have similar properties
• Metals are on the left side and middle of the table.
o (i) Alkali metals are very reactive metals (e.g. Na)
o (ii) Alkaline earth metals somewhat reactive metals (e.g. Ca)
• Non-metals are in the upper right corner
o (iv) Noble gases- very unreactive gaseous non-metals (e.g. Ne)
o (iii) Halogens- very reactive non-metals (e.g. Cl)
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• Metalloids/Semi-metals form a staircase toward the right side.
• Transition metals include the elements from groups 3- 12. (e.g. Fe, Ni, Cu, Ag, Au)
Periodic Table and Ion Formation
• Recall that an atom is electrically neutral because the number of protons equals the number of electrons.
• Most atoms are capable of either losing or gaining electrons.
• Atoms that have gained or lost electrons are electrically charged particles called ions
• When there are fewer electrons than protons, the ion is positively charged
o Metal atoms lose electrons to form positively charged ions called cations
o Many metals form a cation one way. Eg. Al +3
o Some metals are multivalent meaning they can form ions in more than one way. Eg Fe +3, +2
• When there are more electrons than protons, the ion is negatively charged.
o Non metals with very few exceptions, gain electrons to form negative ions called anions
Table 2: Examples of Atoms and Ions
Symbol Charge Atom or Ion # of protons # of electrons
C
Al +3
S-2
Na
Use your periodic table to complete the following:
Symbol Charge Atom or Ion # of protons # of electrons
F
I
Ar
+2 20
0 14
-3 7
1 1
10 10
11 10
17 18
Atom 2
23
+3 81
92
Bhor Diagrams
• Niels Bohr who showed the arrangement of ______________ in energy levels (or shells)
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• Bhor diagrams show how many protons, found in the nucleus and
electrons are found in the shells/ orbits surrounding the nucleus
How to draw Bohr diagrams:
• draw a nucleus and identify how many protons and neutrons there are
• draw electron shells around the nucleus and add the appropriate amount of electrons using 2 rules:
o in neutral atoms (no charge) the number of e- equals the number of p+
o a maximum of 2 can be in the first shell.
o a maximum of 8 can be in shells 2 and 3
Ex.1 Potassium
Atomic number =
Atomic mass =
# of protons=
# of electrons=
# of neutrons=
Ex. 2 Chlorine
Atomic number =
Atomic mass =
# of protons=
# of electrons=
# of neutrons=
1. Draw Bohr diagrams for the first four Alkali Metals (i.e. H, Li, Na, K). What pattern do you observe as you
move from one period to the next?
2. Draw Bohr diagrams for the first four elements in the fourth period (i.e. K, Ca, Sc, Ti). What pattern do you
observe as you move from one family to the next?
3. Draw Bohr diagrams for the first three Noble Gases (ie. He, Ne, Ar). What similarity do you observe
between all three?
Patterns of Electron Arrangement in Periods
• The period number of an element ____equals the number of occupied shells of its atoms
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• For ex., the two elements in period 1, hydrogen and helium have one occupied shell.
• The elements in period 2 have 2 occupied shells and the pattern continues as evident in the figure below.
• As you move from left to right across a period, one more electron is added to the outermost shell (valence
shell)
• Notice that neon and argon have eight electrons in their valence shells. This is referred to as a”Octet
Rule”: Alll atoms ‘want’ to achieve this state by gaining or losing e-
Patterns of Electron Arrangement in Groups
• Valence electrons, electrons in the valence or outer shell, are involved in chemical bonding
• Group 1 elements have _____1___ valence electrons
• Group 2 elements have ____2____ valence electrons
• Group 13 elemetns have ____3____ valence electrons
Use the table above along with your periodic table to answer the following questions:
1) Based on the patterns of the periodic table, identify the number of occupied shells for each of the
following elements:
a. calcium ___4_____ c. sulphur _______2___
b. krypton ____4______ d. iodine ____7______
2) Based on the patterns in the periodic table, identify the number of valence electrons for eachof the
following elements:
a. chlorine_______7______ c. strontium_____2_____
b. magnesium_____2_____ d. bromine_____7______
Forming Compounds
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• When two atoms move closer together, their valence electrons interact.
• A chemical bond forms between atoms if the new arrangement of electrons is stable.
• Lower energy states are more stable. The lowest energy states are achieved when the atoms in the
compound have the same arrangement of valence electrons as the noble gas they are closest to in the
periodic table.
o Referred to as Noble Gas stability. The noble gases are stable (i.e. unreactive) because their valence
shells are completely full
• Atoms acquire stable valence shells, when they form compounds, in one of three ways:
o Metal atoms may lose electrons to other atoms, forming cations (+) .
o Non- metal atoms may gain electrons from other atoms, forming anions (-)
o Atoms may share electrons
o Reactivity of an element is linked to how close the atom is to having a full valence shell
o Therefore we find the most reactive elements are those of Groups 1 and 17
• Compounds are either ionic or covalent
o An ionic compound contains a metal, (+) ion, and a non metal, the (-) ion.
o In an Ionic bond, one or more electrons transfers from each metal atom to each non-metal
o In covalent bonding, the atoms of many non-metals share electrons with other non- metal atoms
o A covalent molecule is a group of atoms in which the atoms are bound together by sharing one or
more pairs of electrons
o The pair of electrons involved in a covalent bond are sometimes called the Bonding Pair
o A pair of electrons in the valence shell that is not used in bonding is called a Lone Pair
o Bohr diagrams, as seen below, can be used to describe simple covalent compounds.
Lewis Diagrams
o Gilbert Lewis (1875-1946), an American chemist, invented a method of showing bonding
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o A Lewis diagram illustrates chemical bonding by showing only an atom’s valence/outer electrons and
chemical symbol
o Lewis diagrams are also referred to as Lewis Dot diagrams or Lewis structures.
o The figure below compares Bohr diagrams to Lewis diagrams:
Rules for drawing Lewis diagrams:
o Dots representing electrons are placed around the element symbol at the points of a compass
(N,E,S,W)
o Electron dots are placed singly until the 5th electron is reached, then they are paired.
Lewis diagrams of ions:
o For positive ions, one electron dot is removed from the valence shell for each positive charge of the
ion. (this usually means all of the electron dots are removed)
o Square brackets [] must be placed around the element symbol with the ion charge shown outside of
the brackets at the top right.
o For negative ions, one electron dot is added to each valence shell for each negative charge of the
ion. (This usually means the element’s symbol is surrounded by eight electron dots)
o Square brackets are also placed around the element symbol with the with the ion charge at the top
right, outside of the brackets. (refer to figure below).
NaCl à
Lewis diagrams of compounds:
o Lewis diagrams can be used to show ionic compounds as seen in the figure below:
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o Lewis diagrams are also used to show covalent compounds as seen in the figure below:
o Covalent molecules can also be represented by Lewis diagrams:
Lewis diagrams are also used to explain why some of the non-metal elements exist as diatomic molecule
A diatomic molecule is a pair of identical atoms that are joined by covalent bonds
Diatomic elements form this way because the two atom molecules are more stable than the individual atoms.
o The diatomic elements include:
o hydrogen: H2
o nitrogen: N2
o oxygen: O2
o chlorine: Cl2
o bromine: Br2
o fluorine: F2
o iodine: I2
4.2 Names and Formulas of Compounds
Ionic Compounds:
Naming Ionic Compounds from their Formula
1. Name the metal ion first
2. Name the non metal second, but change the end to “ide”
3. Combine the two.
4. Ignore the subscripts
Examples:
a. CaI2 b. MgCl2 c. AlCl3 d. NaCl e.KBr
Calcium iodide magnesium chloride aluminum chloride sodium chloride potassium bromide
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Formulas of Ionic Compounds:
1. They symbol from the periodic table is used as is the charge
2. To write the formula you must balance the charges of the (+) and the (-) ions by having a different number
of atoms of each of them
3. There are two methods to do this. By a. balancing charges or b. criss cross
4. You need to indicate the number of atoms of each element in the formula using subscripts. If you have 1
atom of an element you do not need to include the subscript. If you have more than one of an element you
need to include the subscript.
*Note: You must reduce subscripts to lowest terms if possible
Ex. magnesium chloride
Multivalent Metals
o A multivalent metal is one that can have more than one ion charge
Examples: Cu +1 or +2 Pb +2 or +4 Fe +2 or +3
o Naming:
o The charge on the metal is important to its properties, so we must be specific when we write the
name of a compound with one of these metals
o Roman Numerals are used to say which ion charge was used.
§ These are placed between the name of the metal and the non metal for ionic compounds
Ex. Cu2O copper (I) oxide
Writing Formulas:
• Writing formulas for multivalent metals is the same as any other ionic compound, you just need to make
sure you are using the right ion charge .
• Once you have the correct charge use the criss cross method or balancing charges method
o Example: Lead (IV) oxide
Polyatomic Ions:
o A polyatomic ions 2 or more of atoms joined by covalent bonds carrying an overall charge.
o Examples: NH4 +1 or OH-1 or (ClO)-1
**Charge&R.Ntable
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o They act as unit with a charge and don’t get split into their individual atoms
o You will be given a list of these on the back of your periodic table for quizzes and tests.
o All of these are (-) charged, except for NH4 +1 and H30+1
Formulas of Compounds with Polyatomic Ions:
1. Figure out which ion is (+) and which ion is (-).*The first one in the name must be the +ion
2. Use the balancing charges or criss cross method to come up with a formula
3. Use subscripts to show how many of each ion you need by adding it afterwards.
a. If there is more than one polyatomic ion, you must put brackets around it before adding your
subscript on the outside.
Example: magnesium hydroxide
4. You do NOT change the ending of the polyatomic ion. Ex. NaNO3 = sodium nitrate
Naming of Compounds with Polyatomic ions:
Hint: If there is more than two capital letters in the fomula there is a polyatomic ion present in the compound
o Identify the polyatomic ion. Then name the molecule according to the same rules as regular ionic
compounds
o Example Ca(OH)2
Calcium hydroxide
Prefix Number Mono 1 Di 2 Tri 3 Tetra 4 Penta 5 Hexa 6 Septa 7 Octa 8
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Covalent Compounds
Naming Binary Covalent Compounds:
o Remember, covalent means non-metal + non metal
o When you have two non-metals joined together, we
Name them in a specific way
o We use prefixes to help us
1. Name the element closest to the left of the Periodic Table first.
2. Name the other element second and change the ending to “ide”
3. Add the appropriate prefix before the element saying how many of the element are in the compound.
4. If there is one of the first element, you do not write the prefix
5. Do not reduce numbers to lowest terms
Example: N2O2
dinitrogen dioxide
Writing Formulas:
1. Write the element symbols in the order they appear in the name
2. Add subscripts to the symbols based on the prefixes of each element.
Example: Carbon tetrasulfide
CS4
Nona 9 deca 10
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Summary of Naming
Before you start to name a compound, you must ask yourself a couple of questions.
TYPE 1. Is this a simple ionic compound? (metal + non-metal )
*metal stays the same, non metal’s ending changed to –ide
ex. ZnS = zinc sulphide
TYPE 2. Does the metal have more than one option for its charge? (*Check the periodic table )
* You must put Roman numerals ( ex: in brackets between the metal and the non-metal if it does)
Ex. iron (II) oxide = Fe+2 + O2-
FeO
TYPE 3. Is there a polyatomic ion in the compound? (There are more than 2 elements in the compound or there are
brackets in the formula)
*Look on the back of your periodic table to find the name of the polyatomic ion
Ex. lithium acetate= LiCH3COO
TYPE 4. Is it a covalent compound? (Non-metal + non metal)
* You must use prefixes (ex. mono- or di-) in the name.
Ex. CO2 : Carbon dioxide
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4.3 Chemical Equations:
A + B à C
Reactants Products
-both reactants and products are pure substances ( either elements or compounds)
- reactants and products have different properties
- chemical reaction shows chemical changes
- scientists use equations to show how the atoms are rearranging in the reaction
- you can write an equation with names or with chemical symbols
Word Equation: nitrogen monoxide + oxygen à nitrogen dioxide
Symbol Equation: 2NO(g) + O2 (g) à 2NO2 (g)
Writing and Balancing Chemical Equations
• The simplest form of chemical equation is a word equation.
o Potassium metal + oxygen gas à potassium oxide
• A skeleton equation shows the formulas of the elements/compounds.
o A skeleton equation shows atoms, but not quantities of atoms
§ K(s) + O2(g) à KO
• A balanced chemical equation shows all atoms and their quantities
o Balancing makes sure that the number of each atom is the same on both sides of the reaction arrow
o Always use the smallest whole-number ratio
§ 2K + O2 à 2KO
Counting atoms to balance an equation
• Using the law of conservation of mass, we can count the atoms on each side and balance them in
chemical equations
o Word equation: methane + oxygen à carbon dioxide+ hydrogen
o Skeleton equation:
o Balanced equation:
Reactants Products
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o To balance compounds, take note of how many atoms of each element occur on each side of the
reaction equation.
o To balance, change the numbers of molecules involved by adding a coefficient so the atoms on
both sides of the equation are equal (*T table)
Strategic Balancing Tips
o Balance the atoms in compounds first
o Balance one compound at a time
o Polyatomic ions (i.e. NO3– ) can often be balanced as a whole group
o Be careful of special cases such as polyatomic ions and the ‘special seven’: HOFBrINCl
o Several common covalent molecules containing H have common names that you should know
§ Methane= CH4
§ Glucose= C6H12O6
§ Ethane= C2H6
§ Ammonia= NH3
• Only change coefficients, never subscripts!
• If H and O appear in more than one place, always balance them LAST!
• Always double check at the end!
