Chemistry Sections 5.1, 5.2, 5.5, 5.6, 5.7, 5.8, 5.9, 5.11, 5.12, 5.13, 6.5, 6.6, 6.7, 6.19, 7.1, 7.2, 7.3, 7.5,
An introduction to chemistry • Chemistry can be defined as the study of chemicals and their reactions. • Chemicals may be described by their physical characteristics or their chemical characteristics; • Physical characteristics include things like colour, state at room temp., smell, boiling or melting points. • Chemical characteristics mean how a chemical reacts with other chemicals. A chemical change occurs when a substance changes to a new substance.
Mixtures • Most chemicals exist in nature as mixtures, made up of 2 or more substances. • These mixtures may be either homogeneous or heterogeneous. • Homogeneous mixtures are those in which the components are not distinguishable, is completely uniform. Ex coffee or chocolate ice cream • Heterogeneous mixtures are those in which the components are distinguishable. Ex rocky road ice cream, stew
Homogeneous mixtures- • Heterogeneous mixtures-
Pure Substances • Are not as common as mixtures, consist of elements or compounds • Elements are the simplest form of matter that can exist under natural conditions. Ex. Hydrogen, carbon, sodium • Compounds are pure substances that contain two or more different elements in fixed proportions. • Compounds are usually identified with a chemical formula, acombination of letters and numbers to tell you what type and how many of each element is present. Ex. H2O
WHMIS • Stands for Workplace Hazardous Materials Information System • Is a system to inform those using or exposed to chemicals the hazards they may encounter. • Every chemical used in the school (cleaners included!) comes with a MSDS (Material Safety Data Sheet) that describes hazards associated with the chemical, disposal procedures etc. • Complete questions 1, 4, 10 – 12 on page 175
Section 5.5 Elements and the Periodic Table
The periodic table • Organizes elements according to their atomic structure, physical and chemical properties. • The columns (up and down) are known as groups and the rows (across) are known as periods • Chemical families are groups of elements that have similar properties • We can use the organization of the elements in the periodic table to predict their reactivity (how well an element will react)
The Periodic Table • Interactive Periodic Table • Periodic Table: Groups and Trends
Elements An elementis a substance made up of only 1 type of atom. There are about 112 different elements that make up the periodic tableof the elements. On the periodic table each atom type has its information. For example… Atomic no. Symbol Name Mass no.
11 B 5 Periodic Table Atomic no. Any atom can be identified by the atomic no., the symbol or by the name. For instance... Symbol Name H 1 Iron 26 Magnesium Mg The information from the table can also be shown as:
Questions pg 184- 186 • Using table 1 on page 185, compare metals to non metals. • Where can metals and nonmetals be found on the periodic table. • Describe the four chemical families of the periodic table. • Fill in the following table about sub atomic particles
11 B 5 What it means The Atomic Number: = number of protons = number of electrons (as an atom has the same of each) The Mass Number: = number of protons + neutrons - why are electrons not included in the mass no? So for Boron… Protons = Electrons = Neutrons = What about Phosphorus? Protons = Electrons = Neutrons = 5 15 5 15 5.811 16
Electron configuration • Electrons travel in orbits or orbitals around the nucleus. The atomic number on the periodic table tells you how many electrons each element has. • Because atoms are electrically neutral, the number of electrons equals the number of protons.
ELECTRON ARRANGEMENT Electrons are very fast moving. They are arranged in shells around the nucleus. The first shell fits… The second fits… The third fits… So the electron shell for 12Mg would be… 2 e 8 e 8 e 2, 8, 2 Interactive periodic table
Ionic Bonding • Na + Cl 2,8,1 2,8,7
IONIC FORMULAE So Mg2+ will be attracted to Cl-. Because Mg is 2+ and Cl is only 1-, Mg will attract 2 Cl’s. The compound formed will be MgCl2. The subscript shows that there are 2 Cl’s for each Mg. If the starting ions were Cu2+ and S2-, the 2 ions have the same charge. So each Cu will only attract 1 S. The compound formed will be CuS. There are never any charges on the final product - they balance out
6 e 2 e 8 p + 8 n º The mass number tells you the mass of the element and when rounded to the nearest whole number can be used to determine the number of neutrons inside the nucleus. Mass number – atomic number = # of neutrons. Ex. Oxygen Atomic # = 8 Mass # = 16
Ions Elements are most stable where their outer electron shell or orbit is full. Elements whose orbit are almost full lose or gain electrons, and become ions to achieve stability Elements that gain electrons (and therefore a negative charge) form anions. Elements that lose electrons (and therefore have a positive charge) form cations.
Anions • Are formed when non-metals gain electrons. • What was once a neutral atom becomes a negatively charged ion. • The value of the charge is equal to the number of electrons gained.
Cations • Are formed when metals lose electrons • What was once a neutral atom becomes a positively charged ion. • The value of the charge is equal to the number of electrons lost
Naming Ions • Cations are named by simply stating the element from which it forms followed by the word “ion” • Ex. Sodium ion • Anions are named by stating the elements from which it forms and replacing the ending with “ide” • Ex. Chloride
Compounds: Ionic Bonding How do atoms become stable ions? Ionic bonding animation
Types of Ions… • Anions…Number of electrons is greater than the number of protons • Negative charge • Cations…number of electrons is less than the number of protons. • Positive charge
Determining ion…general guidelines • Metals form cations • Non-metals form anions
Writing formulas… • Five step rule… • 1. Write the symbol. • 2.Write the charges. • 3. Cross over the charges from top to bottom. • 4. Remove the charge. • 5. Simplify the numbers. • Formulas...
Predicting Ionic Charges Group 1: Lose 1 electron to form 1+ ions H+ Li+ Na+ K+
Predicting Ionic Charges Group 2: Loses 2 electrons to form 2+ ions Be2+ Mg2+ Ca2+ Ba2+ Sr2+
Predicting Ionic Charges Group 3: Loses 3 electrons to form 3+ ions B3+ Al3+ Ga3+
Predicting Ionic Charges Group 4: Lose or gain 4 electrons? Neither!Group 13 elements rarely form ions.
Predicting Ionic Charges Nitride N3- Group 5: Gain 3 electrons to form 3- ions P3- Phosphide As3- Arsenide
Predicting Ionic Charges Oxide O2- Group 6: gain 2 electrons to Form 2- ions S2- Sulfide Se2- Selenide
Predicting Ionic Charges F1- Fluoride Br1- Bromide Group 7: gain 1 electron to form 1- ion Cl1- Chloride I1- Iodide
Work for today... Chapter 5.5 • 1. Describe the alkali metals. • 2. How are the alkali metals different from the alkali earth metals? • 3. Describe the noble gases. • 4. Describe the halogens. Chapter 5.6 Do #’s 1, 2, 3, 4 Chapter 5.8 Do #’s 1,2,3,4,5,6
Predicting Ionic Charges Many Transition metals have more than one possible ionic charge Iron(II) = Fe2+ Iron(III) = Fe3+
Predicting Ionic Charges Some transition elements have only one possible charge Zinc = Zn2+ Silver = Ag+
Writing Ionic Compound Formulas Example: Barium nitrate 1. Write the formulas for the cation and anion, including CHARGES! ( ) 2. Check to see if charges are balanced. Ba2+ NO3- 2 Not balanced! 3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion.
Writing Ionic Compound Formulas Example: Ammonium sulfate 1. Write the formulas for the cation and anion, including CHARGES! ( ) NH4+ SO42- 2. Check to see if charges are balanced. 2 3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced!
Writing Ionic Compound Formulas Example: Iron(III) chloride 1. Write the formulas for the cation and anion, including CHARGES! Fe3+ Cl- 2. Check to see if charges are balanced. 3 3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced!
Writing Ionic Compound Formulas Example: Aluminum sulfide 1. Write the formulas for the cation and anion, including CHARGES! 2. Check to see if charges are balanced. Al3+ S2- 2 3 3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced!
Writing Ionic Compound Formulas Example: Magnesium carbonate 1. Write the formulas for the cation and anion, including CHARGES! Mg2+ CO32- 2. Check to see if charges are balanced. They are balanced!
Writing Ionic Compound Formulas Example: Zinc hydroxide 1. Write the formulas for the cation and anion, including CHARGES! ( ) 2. Check to see if charges are balanced. Zn2+ OH- 2 3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced!
Writing Ionic Compound Formulas Example: Aluminum phosphate 1. Write the formulas for the cation and anion, including CHARGES! 2. Check to see if charges are balanced. Al3+ PO43- They ARE balanced!
Naming Ionic Compounds(continued) Metals with multiple oxidation states • - some metals form more than one cation • - use Roman numeral in name • PbCl2 • Pb2+is cation • PbCl2 = lead(II) chloride • Roman numeral is equal to the charge of the cation
Complex Ions (Polyatomic Ions) • Mg 2+, I-, Li +, S2- • are all called simple ions or monatomic ions • Complex, or polyatomic ions, are tightly bound groups of ions that behave as a unit and carry a charge. Example : sulfate ion. A sulfate ion is composed of 1 sulfur atom and 4 oxygen atoms. These 5 atoms together form a unit with a charge. • SO4 2- • Recognizing complex ions is a key in naming chemical compounds and writing chemical formulas. Polyatomic Ion Rap...
Ammonium……………... Nitrate…………………… Permanganate…………. . Chlorate………………… Hydroxide………………. Cyanide…………………. Sulfate…………………... Carbonate………………. Chromate……………….. Acetate………………….. Phosphate………………. Polyatomic Ions • NH4+ • NO3- • MnO4- • ClO3- • OH- • CN- • SO4 2 - • CO32- • CrO42- • C2H3O2- • PO43-
cobalt (III) carbonate • Co2(CO3)3 • beryllium nitrate • Be(NO3)2 • Polyatomic tutorial... • Please do #’s 1,2,3,4,6,7 on page 189 Chapter 5.9
Molecular Compounds Section 5.11