Copy of Chemistry Notes 2024-2025 (Hillaby).gslides

Science 10 Lab Safety Properties of Matter Atomic Theory Lab Safety Before Going Into the Lab Checklist 1. Wear lab goggles. 2. Ensure long hair is tied up. 3. Roll up any loose or long sleeves. 4. Shoes must cover up the entire foot. No open toed shoes. 5. Wear a lab apron/lab coat. 6. Read the lab ahead of time. 7. Ask questions if you are confused First Aid If your skin contacts a chemical, rinse for 15 mins under cool water If you get anything in your eyes, rinse for 15 minutes If you burn yourself, rinse under cool water for 10-20 mins Always let a teacher know if a spill, breakage, or accident occurs, regardless of how small ○ Certain chemicals require special cleanup techniques After Lab Dispose of all materials as directed ○ Some chemicals can be poured down the sink with lots of water ○ Others will require special disposal, and containers/directions will be provided Clean all equipment, and your work area Wash your hands WHMIS Stands for Workplace Hazardous Materials Information System WHMIS is a system by which we can inform people in a workplace, such as a lab, the hazards that certain chemicals possess These symbols tell us what the danger is so that we know what personal protective equipment to wear and how to treat the chemicals WHMIS just had a change in 2015. WHMIS Symbols Symbol Practice A. B. C.. D. SDS (Material Safety Data Sheet) In our schools, these chemicals will also come with an information sheet, known as a Safety Data Sheet , which contains the following information: 1. Chemical and Physical Hazards 2. Toxicity and Health Eﬀects 3. First Aid 4. Spill and Leak Cleanup Procedures Chapter 1: Atoms, Elements and Compounds Describing Matter → Review of Science 9 Properties of Matter Physical Properties Properties of a substance that relate to the kinds of physical changes that take place Physical Change - When the shape or state is altered, but the substance stays the same. Physical Properties Examples Boiling Point - temperature at which a substance boils Malleability - ability to be beaten or rolled into sheets without crumbling Ductility - ability to be stretched without breaking Solubility - ability to dissolve Crystal Formation - crystalline appearance Conductivity - ability to conduct heat or electricity States of Matter Matter: Anything that has MASS and TAKES UP SPACE Matter can be found in three diﬀerent states: Gases: NO deﬁnite shape, NO deﬁnite volume Liquids: NO deﬁnite shape but deﬁnite volume Solids: Deﬁnite shape and deﬁnite volume Practice: Other States of Matter SUPER HOT Discovered in 1879 Found in stars, northern lights, ﬂuorescent light bulbs and neon signs SUPER COLD (- 273 degrees) Found in 1955 Atoms clump together to form a blob Chemical Properties How a substance interacts with other substances. Chemical change- occurs when a substance changes into a diﬀerent substance Chemical Properties Examples Ability to burn - combustion Behaviour in air - tendency to degrade, react, or tarnish Reaction with water - tendency to corrode Reaction with acids - corrosion, sometimes bubble formation Reaction to heat - tendency decompose Reaction to red and blue litmus - red=acid, blue=base, no colour change=neutral Classifying Matter The Atomic Theory Numerous scientists have tried to explain the STRUCTURE and BEHAVIOURS of ATOMS The atomic theory is STILL EVOLVING TODAY! Each picture illustrates a scientists theory of what the “atom” looked it ! 1. John DALTON (1808) ALL MATTER is made up of small particles called ATOMS ○ The ATOM is a solid uniform sphere Atoms cannot be CREATED or DESTROYED or DIVIDED (we now know this to be true) All atoms of the same element have the SAME MASS and SIZE (not true) COMBINING different elements forms COMPOUNDS Chemical reactions change how atoms are GROUPED → The atoms themselves are not changed in the reaction Also known as the “Billiard Ball Model” 2. J.J. Thomson (1902) Atoms COULD be divided Atoms contain tiny, negatively charged particles called ELECTRONS These tiny electrons float in a spherical cloud of positive charge Known as the “plum pudding” or “raisin bun” model 3. Rutherford (1909) Made a “SOLAR SYSTEM” atomic model Most of the atom is made up of empty space where the ELECTRONS “orbit” around the nucleus The NUCLEUS of the atom contains: 1. Protons → large POSITIVELY charged particles 2. Neutrons → large NEUTRAL particles Rutherford Experiment! Scattering experiment (EVIDENCE) he used to prove there was a NUCLEUS Evidence that if an atom’s structure was like pudding (as Thomson said) + particles beamed at it should just go straight through They didn’t! They got deflected by something + in the nucleus (charges repelling each other) 4. Bohr (1913) Refined Rutherford’s Model! ELECTRONS don’t “orbit” because their energy would cause them to crash into the nucleus if they did! Electrons travel only in specific ENERGY LEVELS or ELECTRON SHELLS Electron SHELLS → they hold a speciﬁc amount of electrons! 5. Current Atomic Theories Energy levels are divided into SUBLEVELS or electron “clouds” → NOT in a FIXED ORBIT Electrons travel in waves, which means their exact positions cannot be determined Since they can be anywhere, scientists can only make educated guesses as to the positions of the electrons around the nucleus Neutrons and Protons are made of even SMALLER particles called QUARKS the Atomic Theory keeps evolving as more experiments are done! More ﬁndings! Current Model: Electron clouds, etc Clearly Stated: What is an Atom? The 2,400-year search for the atom - Theresa Doud Science 10 Day 2: Periodic Table and Atoms Intro Elements and the Periodic Table The last four elements (in yellow) are the last 4 synthetic elements found in 2016! Dmitri Mendeleev (1834-1907) Dmitri Mendeleev made the periodic table in 1869, which has been continuously improved upon. The periodic table has many pieces of information about an element Periodic Table of Elements Elements are arranged on the periodic table according to 4 BASIC PATTERNS 1. Atomic Number 2. Metal/ Non-metal properties 3. Period 4. Groups 1. Atomic Number Represents the number of PROTONS for that element NOTE: This is also the number of ELECTRONS for an ATOM of that element 1 11 15 H Na P Note: The # of electrons will not always be the # of protons when we deal with IONS Getting to KNOW the elements on the periodic table Find the following information for Nitrogen Atomic Symbol Atomic Number Ion Charge Group Number Period Number Find the following information for Calcium Atomic Symbol Atomic Number Ion Charge Group Number Period Number 2. Metals vs. Non-Metal Properties The periodic table is divided into metals and non-metals Metals on the left of the “stairs” and non-metals on the right There is a “staircase” along the staircase is metalloids (they have properties of metals and non-metals 3. Periods Periods are the horizontal rows on the periodic table There are 7 periods Each period has the SAME NUMBER OF ELECTRON ENERGY LEVELS Example: The period 4 atoms each have 4 energy levels 4th Energy Level K (Potassium Atom) Fe (Iron) Atom Kr (Krypton Atom) 4. Groups Groups are the VERTICAL columns on the periodic table There are 18 groups Groups have similar REACTIVE PROPERTIES Group Names/ Properties Group 1: Alkali metals (HIGHLY reactive), shiny → (H+ excluded → its a non-metal, though very reactive) Group 2: Alkali earth metals (light, reactive), form oxide coating when exposed to air Group 3-12: Transitional metals Group 18: Noble gases (stable non-metals) Reactivity Reactivity of metals increases as you go down and to the left ○ Therefore, francium is the most reactive metal. Noble gases are inert (does not react with others). Reactivity of non-metals increases as you go up and to the right ○ Therefore, ﬂuorine is the most reactive non-metal 1. BLANK periodic table 2. Colour: Metals/ Nonmetals/ Metalloids 3. Label: Groups/ Periods Comparing Metals and Non- Metals METALS NON-METALS Properties of Metals Metals are SOLID at room temperature with the exception of mercury which is a liquid Metals have a SHINY LUSTRE (think of the metal on a brand new car) Metals are good CONDUCTORS OF HEAT AND ELECTRICITY Metals are MALLEABLE ( they can form a thin sheet by hammering it or rolling it Metals are DUCTILE (they can be stretched into wire) Properties of Non-Metals Non-metals can be solid, liquid or gas Non-metals are not very shiny Non-metals are POOR CONDUCTORS OF HEAT AND ELECTRICITY Non-metals are BRITTLE (they break apart if hammered) Non-metals are NOT DUCTILE Properties of Metalloids Metalloids (staircase) that separates metals from non-metals, they have properties of both All metalloids are solid at room temperature Some metalloids are shiny while others are dull Some metalloids conduct ELECTRICITY but none are good conductors of heat All metalloids are brittle and non-ductile The Atom Atoms have the same number of protons and electrons The nucleus is made up of neutrons and protons. Net charge = 0 ○ Neutrons = no charge ○ Protons = + 1 charge ○ Electrons = - 1 charge But! the NUCLEUS has a positive charge, and is most of the mass of the atom Calculating the Average Number of Neutrons in an Atom The atomic mass of an atom is calculated by adding together the # of protons and the # of neutrons (electrons are too small to worry about) The neutron # is not always the same, the atomic mass on the periodic table reflects the most common # of neutrons found in the element added to the protons # of Neutrons = Atomic Mass - Atomic # Roundest to the Number of Protons nearest whole number Ions Created from a neutral atom losing electron(s) or gaining electron(s) Number of protons DOES NOT equal to number of electrons Either ends in -ide or has the word “ion” Net charge is positive or negative Cations – positively charged ions Anions – negatively charged ions Metals tend to lose electrons (CATIONS) Non-Metals tend to gain electrons (ANIONS) Atomic Structure Figuring Out Protons, Neutrons, and Electrons Atoms Ions # of Protons = Atomic Number # of Protons = Atomic Number # of Electrons = # of Protons # of Electrons = Atomic Number – Remember: Atoms have a Ion Charge (make sure you are 0 charge. keeping the - charge) # of Neutrons = Atomic Mass - # # of Neutrons = Atomic Mass - # of of protons (round to the nearest protons #) Example- Fluorine Fluorine (Note: Fluoride is the ion. Fluorine is the neutral atom) Atomic Number: 9 Mass Number: 19.00 # of protons = atomic number = 9 # of neutrons = mass number - atomic number = 10 # of electrons = # of protons = 9 (because fluorine is a neutral charge atom) Example- Fluoride ANSWERS Fluorine (Note: Fluoride is the ion. Fluorine is the neutral atom) Atomic Number: 9 Mass Number: 19.00 # of protons = atomic number = 9 # of neutrons = mass number - atomic number = 10 # of electrons = Atomic number - ion charge = 9 - (-1) = 9+1 = 10 Calculate the Most Common Number of Neutrons in these Atoms REMEMBER Atomic Mass - Atomic # = # of A. Lithium neutrons B. Nitrogen C. Potassium D. Silver E. Sulfur ANSWERS REMEMBER Atomic Mass - Atomic # = # of A. Lithium 7 - 3 = 4 neutrons neutrons B. Nitrogen 14 - 7 = 7 neutrons C. Potassium 39 - 19 = 20 neutrons D. Silver 108 - 47 = 61 neutrons E. Sulfur 32 - 16 = 16 neutrons More Practice Element Protons Neutrons Electrons Nitrogen 49 49 Sodium 6 Check Answers Element Protons Neutrons Electrons Nitrogen 7 14-7 = 7 7 Indium 49 115-49 = 66 49 Sodium 11 23-11 = 12 11 Carbon 6 12-6 = 6 6 Science 10 Atoms and Ions Atoms have NO net charge - REVIEW In an atom the # of positive protons is always the same as the # of negative neutrons so the net charge is ALWAYS ZERO The atomic # tells you how many protons and electrons are in the atom. Remember for IONS the electrons will not be the same # as the protons Example - Carbon Atom The Atomic # of CARBON is 6 There are 6 protons (+), 6 neutrons (0) (most commonly) and 6 electrons (-) Because the 6+ balances the 6- there is NO CHARGE IN THIS ATOM Energy Levels → Based on Bohr’s Model There are different energy levels on the electron shells around the nucleus that can fit certain amounts of electrons ○ The FIRST shell can fit 2 ○ The SECOND shell can fit 8 ○ The THIRD shell can fit 8 Note: The number of elements in the Period = the max number of electrons in each energy level! Energy Levels How many VALENCE electrons does this Electrons are drawn filling the innermost levels first atom have? Valence Electrons: any electrons in the outermost shell these are involved in chemical reactions can be shared or transferred Note: The Last Digit of the Group Number is the number of valence electrons! Neat! Drawing a Bohr Model The nucleus contains the protons and neutrons The electrons fill the energy levels starting with the closest shell to the nucleus until all the electrons have been used up Ex. Oxygen (Atomic #8) There are 8 Protons and 8 Neutrons There are 8 electrons - 2 in the ﬁrst shell - 6 in the second shell Draw Bohr Models! 1. Neon (Atomic # 10) Draw Bohr Models! 2. Sodium (Atomic # 11) Draw Bohr Models! 3. Chlorine (Atomic # 17) Energy Level - Sodium (Na) Another way to draw the atom is a simple energy level diagram which saves you from having to draw in all the electrons. Step 1: Determine the # of protons, neutrons, electrons Step 2: Show protons and neutrons in the nucleus by writing the # of each and drawing a circle around them Step 3: Add electrons to each ring (starting closest to the nucleus) until all the electrons have been placed. It is easier to write the electrons as numbers than dots. indicate the level using a curved line. Energy Level Diagrams Practice Draw the energy level diagrams for the following atoms 1. Lithium 2. Chlorine 3. Carbon Draw energy level diagrams for the ﬁrst 20 elements. What patterns do you notice? Helpful Hints- Energy Diagrams The PERIOD # = # OF ENERGY LEVELS The GROUP # = # OF ELECTRONS IN THE LAST ENERGY LEVEL ○ Note: Electrons in the last energy level are called VALENCE ELECTRONS ○ Ex. non-metal Cl: group 17 = 7 valence electrons ○ Ex. metal Ca: group 2 = 2 valence electrons AP ONLY! The Lewis model of the atom is a supermodel ⚫ Lewis Dot Diagrams (LDD) ⚫ these diagrams are a simple model of the electron arrangement in an element or compound ⚫ they represent molecules or atoms joined together by using dots to represent valence electrons ⚫ by pairing dots, bonds are formed until each atom in the molecule gains a complete octet 80 ⚫ LDD Rules: 1. Determine the number of valence electrons of an element ( the number of electron in the last energy level of an atom): 2. Write the element symbol- this represents the nucleus of the element and all its filled orbitals 3. Use one dot to represent each valence electron 81 2. Each side of the element symbol represents an orbital a) each orbital can carry a maximum of 2 e- (2 electrons sharing the same region of space) b) elements have 4 valence orbitals , the maximum number of electrons: 8 ⚫ place one dot into each orbitals before they are doubled up in any orbital ⚫ electrons will occupy all empty valence orbitals before any become full ⚫ Recall: Lone pair -two electrons that occupy the same valence orbital in the atom( they cannot form bonds) 82 Example 1: ⚫ Example: chlorine 7e- 8e- 2e- (17p+) ⚫ Chlorine has 7 valence electrons, therefore ⚫ chlorine Lewis Dot diagram is represented as follows: -each dot represents a valence electron 83 ⚫ Example 2: Sodium ‘s Lewis Dot Diagram ⚫ Example: sodium 1e- 8e- 2e (11p+) ⚫ Na has 1 valence electrons ⚫ there are no lone pairs ⚫ there is 1 bonding electrons ⚫ Na wants to lose one electron ⚫ (you do not have to draw out the electron energy diagrams [Bohr diagrams] each time, as long as you determine the correct number of valence electrons 84 ⚫ Example #3: Draw the Lewis Dot Diagrams atom # of valence Bonding Lone pairs Lewis Dot electrons electrons Diagram carbon nitrogen hydrogen neon 85 Ions! - Why do atoms become Ions? All atoms want to have a stable (full) outer energy level To become stable the atoms will GAIN or LOSE valence electrons (whichever is easier) to ensure their outer energy level is full These electrons are transferred to other ions → THIS IS HOW COMPOUNDS ARE FORMED CAT-ion is PAWistive Positive Ions All positively charged ions are called CATION When an ion LOSES electrons it becomes POSITIVELY charged All METALS are positively charged Negative Ions When an ion GAINS electrons it becomes NEGATIVELY charged ALL NON-METALS are negatively charged All negatively charged ions are called An ANION is a NEGATIVE → ANIONS Like and angry Anaconda Naming Ions To name a METAL ○ Call it the same name and add the word “ion” after it ○ Ex. Sodium Atom → Sodium Ion / Iron Atom → Iron Ion To name a NON-METAL ○ Change the ending to “ide” ○ Ex. Chlorine atom → Chloride/ Bromine atom → Bromide Energy Level Diagram for Ions - Examples A. Beryllium Ion B. Sulfide Common Ion Charges Group 1: 1+ (lose their 1 valence e-) Group 2: 2+ (lose their 2 valence e-) Group 16: 2- (non-metals, gain 2 e-) Group 17: 1- (non-metals, gain 1 e-) Other groups/elements? See “most common ion charge” on periodic table! Note: LEO the lion says GER Losing Electrons is called Oxidation, Gaining Electrons is called Draw energy level diagrams for the ﬁrst 20 ions. Why are some boxes crossed out? What belongs in the box with the purple outline? Science 10 Binary Compounds Chapter 2: Names, Formulas and Properties Binary Compounds Binary Compounds are compounds made up of TWO ELEMENTS Almost always end in “-ide” There are THREE types of binary compounds to know: 1. Ionic Compounds → (metal and non- metal) 2. Molecular Compounds (covalent) → (non-metal and nonmetal) 3. Alloy Compounds → (metal and metal) IUPAC The International Union of Pure and Applied Chemistry is Nomenclature (IUPAC) → started in 1919 by chemists Wanting to STANDARDIZE the way chemists communicate Anytime something refers to the IUPAC way of writing something in chemistry this means it is the universally accepted way of writing something Ionic Compounds CATIONS need to LOSE electrons 2. Ionic Compounds ANIONS need to GAIN electrons A Metal and Non- Metal are attracted to each other → They want to become stable! Metals (+ cations) TRANSFER their VALENCE ELECTRONS to non-metals (- anions) → This is called an IONIC BOND The ratio of (+) charges to (-) charges in an ionic compound must be equal so that the net charges of the molecules = 0 Ionic Bonding Characteristics All IONIC COMPOUNDS share…. 1. They form between METALLIC and NON- METALLIC ELEMENTS 2. They produce IONIC BONDS → this is a strong bond! 3. They involve a CHANGE IN ENERGY → (+) and (-) charges must balance to 0 Ionic Compound Properties Ionic compounds form repeating pattern shapes called CRYSTAL LATTICES Ionic compounds have HIGH melting point → Strong BOND! Hard to break They DO NOT conduct electricity as SOLIDS They DO conduct electricity when melted or dissolved in water (aqueous) → They are called ELECTROLYTES Crystal Lattice - NaCl How to Write the Chemical Formula of an Ionic Compound 1. The symbol of the metallic ion appears FIRST 2. The symbol of the non- metallic ion appears LAST 3. Subscripts (# the comes after the symbol and below) indicate the ratio of ions in the compounds. If NO subscript appears after a symbols you may assume that it is one META L NON-METAL SUBSCRIPT Examples How many of each Na+ and Cl- ion do we need to be Final Product balanced? NaCl + - Each Na+ comes Each Cl- comes in in a package of 1+ a package of 1- We need one Na+ We need one Cl- How many of each ion do we need to be Fe3+ and S2- balanced? + ++ + ++ -- -- -- Each Fe3+ comes in Each S2- comes in Final a package of 3+ a package of 2- Product Fe2S3 We need two Fe3+ We need three S2- Examples Ca2+ and Cl- Examples Ag+ and O2- Rules for Naming Ionic Compounds The name includes both elements in the compound The name of the metallic ion appears FIRST The name of the non-metallic ion appears SECOND → the ending is changed to “-ide” The name of the compound DOES NOT mention the number of ions of each element present in the compound The name of the compound is NOT CAPITALIZED Naming Practice Ionic Compounds CaCl2 NaCl AgI KBr Polyatomic Ions Ions made up of a group of atoms acting as ONE UNIT → “aka blob ions” These atoms are held together with COVALENT BONDS → they have a single charge! See “table of polyatomic ions” on periodic table When you have more than one polyatomic ion in your ionic compound formula use BRACKETS Al 2(SO 4) 3 When you name a compound with a polyatomic ion simply name the METAL first and then the polyatomic ion silver sulfate Examples How many of each K+ and NO3- ion do we need to be Final Product balanced? KNO3 + - Each K+ comes in Each NO3- comes a package of 1+ in a package of 1- We need one K+ We need one NO3- Examples Mg2+ and PO43- Transitional Metals (Multivalent Metals) Transitional metals DO NOT follow the rules that we learned about energy levels→ They can fill outer levels without having the middle levels full?!?! Transitional metals can have more than one charge → Given to you on the periodic table. TWO possible charges! Naming Compounds with Multivalent Metals Scientists use ROMAN NUMERALS in brackets in the chemical name to indicate which CHARGE OF AN ION has been to form the compound Called the STOCK SYSTEM Ex. iron (Fe) can be either: iron(III) Fe 3+ (most common ion charge) OR iron(II) Fe 2+ (other ion charge) Calculating Which ION Charge Has Been Used When examining the chemical compound containing a transition metal with more than one possible charge you can calculate which charge can be used by…. 1. Calculate the net negative charge from the non-metal 2. Examine what charge of the transitional metal would balance the compound Example: What should we call this compound? FeBr2 We must first figure out the net charge of this compound ○ Br- has a charge of -1 and since there are two of them our net negative charge is 2- To name this compound we must determine what charge of iron was used ○ According to the periodic table iron can be Fe 2+ of Fe3+ We must then figure out which charge of iron would balance out 2- ○ Since there is only ONE iron it must contain the charge of 2+ We would therefore write the name of the compound with the charge of the iron (2+) indicated in roman numerals This compound is therefore called iron (II) bromide Name the Transitional Metal MnO2 Ni S 2 3 Cr O 2 3 Au PO 3 4 Write the formula for the Following Transitional Metal chromium (III) oxide platinum (II) bromide palladium (III) iodide gold (III) carbonate Molecular Compounds (a.k.a) Covalent Compounds Molecular Compounds Explained! 1. Molecular Compounds Formed when non-metals SHARE ELECTRONS in their outer energy level to become stable Also known as COVALENT COMPOUNDS The bonds formed when electrons are shared are called COVALENT BONDS Bond is NOT AS STRONG as a IONIC BOND Many of these compounds are ORGANIC COMPOUNDS → Meaning they contain carbon Atoms SHARE electrons so they can both ﬁll their energy levels! Properties of Molecular Compounds Molecular compounds: ○ Can be solids, liquids, or gases at room temperature ○ Have Low Boiling Points ○ Have Low Melting Points → Do not hold their shape when heated ○ Are Electrically Neutral → Do not conduct electricity in solid or liquid form → Called NON-ELECTROLYTES How to Name Molecular (Covalent) Molecules 1. Write the entire name of the FIRST ELEMENT 2. CHange the ending of the name of the SECOND (or last if more than 2) to “-ide” 3. Use a PREFIX to indicate the number of each type of atom in the formula 4. Write the names in LOWERCASE LETTERS S 3Cl 6 = trisulfur hexachloride Prefixes for Molecular Compounds 1. mono – 6. hexa – 2. di – 7. hepta – 3. tri – 8. octa – 4. tetra – 9. ennea – 5. penta – 10. deca - Examples C 6H 9 C 2Cl 4 CO Br 2Cl 3 How to Write the Chemical Formula of a Molecular Compound 1. Write the SYMBOL for the elements in the same order as they appear in the name 2. Use subscripts (# the comes after the symbol and below) to indicate the numbers of each atom of that element. When only one atom is present, omit the subscript C 2Cl 4 Br 2Cl 3 Examples tricarbon tetraoxide silicon dioxide tetraphorsphorous decaoxide triiodine hexabromide Common Names of Molecular Compounds to MEMORIZE These molecular compounds have common names and are not named per usual naming conventions Name Chemical Formula Name Chemical Formula Ethane C2H6(g) Ammonia NH3(g) Propane C3H8(g) Glucose C6H12O6(s) Methanol CH3OH(l) Hydrogen H2O2(l) Peroxide Ethanol C2H5OH(l) Sucrose C12H22O11(s) Ozone O3(g) Methane CH4(g) Water H2O(l) Molecular Elements Atoms of some non-metals form bonds and exist as molecules when NOT JOINED with other elements MEMORIZE (Or remember to check the data booklet!!!!) the list below of elements that form molecules Science 10 Acids and Bases and Solubility Solubility Solubility One way of classifying Ionic Compounds is to say whether or not they are soluble in water ❏ IF the compound is soluble we say it is AQUEOUS (aq) ❏ IF the compound forms a precipitate in water we say it is SOLID (s) ❏ This information should be written at the end as subscript There is a chart (on back of periodic table) to help classify the solubility of compounds Solubility of NaOH (aq) Selected Ionic Compounds PbCl 2 (s) Solubility Rules 1. Locate the negative ion (ANION) on the “Solubility of Selected Ionic Compounds” section on the back of the Periodic Table 2. Underneath find the positive ion (CATION) 3. If the cation is found in the group “ Very Soluble” (High Solubility) it will dissolve easily as is AQUEOUS 4. If the cation is found in the group “ Slightly Soluble” (Low Solubility) it will NOT dissolve easily and will form a precipitate (SOLID) CaCO 3 KBr (NH 4) 2SO 4 FYI- Additional Rules (you will just need to know your solubility table though) Identify if the following compounds are (aq) or (s) NaCl Ni 2S 3 CuSO 4 Au(OH) 3 Properties of Acids and Bases In 1884 Svanté Arrhenius proposed a theory of acids and bases based on their BEHAVIOUR IN WATER ACIDS: dissolve in water to release HYDROGEN IONS H +(aq) BASES: dissolve in water to release HYDROXIDE IONS OH- (aq) BOTH acids and bases are ELECTROLYTES (meaning they conduct electricity) Acids and Bases pH Scale pH scale shows how acidic or basic a substance is! Ranges from 0 - 14 7 is NEUTRAL < 7 = ACIDIC >7 = BASIC What does pH stand for? “Potential for hydrogen (H+) - Higher the # means smaller amount of H+ Acid and Base Indicators 1. Litmus Paper Acid = blue litmus paper will turn RED Base = red litmus paper will turn BLUE 2. pH Paper Acid = result will be BELOW 7 on pH scale Base = result will be ABOVE 7 on pH scale Acid and Base Indicators 3. Phenolphthalein Acid and Neutrals = NO CHANGE/ COLORLESS Bases: turns PINK 4. Bromothymol Blue Acids = turn YELLOW Bases= NO CHANGE/ BLUE Comparing Acids and Bases NAMING ACIDS!!!!! Since acids usually are in a solution when used in reactions, the acid formula should ALWAYS be followed by the subscript (aq) meaning AQUEOUS AQUEOUS means it can be dissolved in water This distinguishes it from a regular ionic compound Any compound that has hydrogen (H+) for it cation and the (aq) symbol at the end should be named an ACID NAMING AN ACID → first figure out the ionic name and then follow these three rules depending on the ending of the anion Rule # 1: If the ionic name ends in “ide” it becomes “ hydro ____________ic acid” Example: HCl (aq) IUPAC name: aqueous hydrogen chloride Acid Name: hydrochloric acid IUPAC : universally accepted way of writing a compound in Rule # 2: If the ionic name ends in “ate” it becomes “ __________________ic acid” Example: HNO 3(aq) IUPAC name: aqueous hydrogen nitrate Acid Name: nitric acid Rule # 3: If the ionic name ends in “ite” it becomes “ _________________ous acid” Example: HNO 2(aq) IUPAC NAME: aqueous hydrogen nitrite Acid Name: nitrous acid The rules for NAMING ACIDS is on your periodic table! Science 10 Acid and Bases /Water Properties The Unique Properties of Water - why is water weird? ONLY water exists in THREE states naturally Diagram of a WATER MOLECULE Water is a MOLECULAR compound The H + and O 2- form COVALENT BONDS (sharing of electrons) Water is a POLAR MOLECULE ….. MEANING there is a (+) and a (-) end of the molecule The (+) of one water molecule gets ATTRACTED to the (-) of another water molecule Forms... HYDROGEN BONDS Unique Properties of Water 1. Water has a HIGH MELTING (0oC) and BOILING POINTS (100oC) This is because its hard to break the HYDROGEN BONDS between the molecules a H IGH t e r has AT Wa IFIC HE 2. It takes a LOT OF ENERGY to increase SPEC CITY A the temperature of water by one degree CAP This is why our bodies don’t change temperature easily! Also why it takes SO LONG to boil water! Unique Properties of Water 3. Water has a CONCAVE MENISCUS and shows CAPILLARY ACTION in containers Water is STRONGLY attracted to the sides of a containers Cohesion: Water molecules tend to stick to other water molecules Adhesion: Water molecules ALSO tend to stick to other things This is how water travels UP the roots of a tree to the top of the leaves! Unique Properties of Water 4. Solid water is LESS DENSE than liquid water This is why ICE FLOATS!!! 5. Water has a HIGH SURFACE TENSION The hydrogen bonds in the water molecule keep it pulled tightly together This is why some bugs can walk on water! Science 10 Day 7: Chemical Reactions and Balancing Equations CHAPTER 3: CHEMICAL REACTIONS Chemical Reactions A CHEMICAL REACTION occurs when one or more substances CHANGE to form different substances (a.k.a a CHEMICAL CHANGE) The substances formed are called the PRODUCTS The substances that undergo the reaction are called the REACTANTS Car rusting formula iron + water + oxygen → rust Reactants Products How Can you Tell a Chemical Reaction Has Occurred? To know that a chemical reaction has occurred there must be a change in ENERGY ❏ Temp change ❏ Emission of light or sound ❏ Electrical energy Other signs of chemical reactions are ❏ Odor of color change ❏ Formation of a gas or precipitate Common Examples of Chemical Reactions How Can You Tell The Difference Between A Physical Change And A Chemical Change? A physical change DOES NOT CHANGE the arrangement of the atoms in a molecule ❏ No energy is taken in or given off ❏ Phase changes are only physical Water heating and cooling changes involves a PHASE ❏ Ex. If you rip a piece of paper it is CHANGE still a piece of paper A chemical change CHANGES the arrangement of the atoms in a molecule ❏ Ex. Burning a piece of paper (heat and odor is released) Indicate if the following are chemical or physical changes? 1. A pellet of sodium 2. Water is heated and turns into steam 3. Sugar dissolves in water 4. A tire is inflated with air 5. A chemical that is heated turns blue Law of Conservation of Energy Law states that energy can be CONVERTED from one form to another but the TOTAL ENERGY REMAINS CONSTANT If other words energy cannot be created on destroyed (it only changes from one form to another) Energy is always required to break chemical bonds and released when new bonds form Exothermic Reactions Reactions that RELEASE energy when new bonds are formed LESS energy is required to break bonds then is released when new bond form EXAMPLES: Explosions, Cellular Respiration, Combustion of Gasoline Endothermic Reactions Reactions that ABSORB energy when chemical bonds are broken More energy is required to break bonds than is released when new bond form EXAMPLES: Cooking Food, Photosynthesis MORE ENERGY LESS ENERGY NEEDED NEEDED Balancing Equations! Why Must Equations Be Balanced? The balancing of equations is required to obey the LAW OF CONSERVATION OF MASS ❖ In the system undergoing change the number of atoms MUST BE CONSTANT ❖ MASS OF REACTANTS = MASS OF PRODUCTS Coefficients A COEFFICIENT is a multiplier for every subscript in a chemical formula it defines (4 H’s and 2 0’s) 2H20 Coeﬃcient Subscript Examples multiplying coefficients- Polyatomic “blobs” A. 3 (PO4) B. 2 (SO3)2 EXAMPLE! This example shows oxygen decomposing into hydrogen gas and oxygen gas: H20 (l) → H2 (g) + 02 (g) Remember the number of each type of atom must be the same on both sides of the equations ➔ 2 hydrogens on the left, 2 hydrogens on the right AWESOME! ➔ 1 oxygen on the left, 2 oxygens on the right UH OH! USE COEFFICIENTS AS MULTIPLIERS! To balance this equation, the number of each type of atom must be the same on either side of the equation… therefore the balanced equation is.. 2H20 (l) → 2H2 (g) + 02 (g) Reactants Products 4 HYDROGENS 4 HYDROGENS 2 OXYGENS 2 OXYGENS Steps to Balancing Chemical Equations 1. Write the chemical formula for each reactant and product, including the state of matter (s, l, g, aq) for each one. This creates the SKELETON EQUATION 2. Try balancing the atom or ion present in the GREATEST NUMBER. Find the lowest common denominator to obtain coefficients to balance this particular atom or ion 3. Repeat step 2 to balance each of the remaining atoms and ions. 4. Check the final reaction equation to ensure that all atoms and ions are balanced! Helpful Hints! 1. Balance METALS first 2. Balance NONMETALS second 3. Balance Hydrogen 4. Balance Oxygen Last Examples! 1. The decomposition of copper (II) oxide into copper and oxygen: Skeleton Equation: ____ CuO(s)→ _____Cu(s) + ______O2 (g) Examples! 2. Sodium and water react to produce sodium hydroxide and hydrogen gas Skeleton Equation: ____ Na(s) +_____H2O(l) ---> ___NaOH (aq)+ ______H2 (g) Examples! 2. Copper and silver nitrate react to produce copper (II) nitrate and silver Skeleton Equation: ____ Cu(s) +_____Ag(NO3)2 (aq) ---> ___Cu(NO3) (s)+ ______Ag (s) Types of Chemical Reactions Types of Chemical Reactions H2 and O2 are 1. Formation Reactions diatomic molecules. Look at your A.k.a synthesis or composition reactions periodic table! Simple elements come together to form compounds Example: 2H2 (g) + O2 (g) → 2H2O (l) Types of Chemical Reactions 2. Decomposition Reactions The products produced are simpler than the reactants. Usually it consists of a compound breaking down into its elements 2H2O (l) → + 2H2 (g) + O2 (g) Types of Chemical Reactions 3. Combustion Reactions Combustion means the ability to burn Always have HYDROCARBONS (CXHX) and OXYGEN as reactants Always have carbon dioxide (CO2) and water (H2O) as products Always have carbon dioxide (CO2) and water (H2O) as products Can be a formation or decomposition reaction! Types of Chemical Reactions 4. Single Replacement Reactions One element is substituted for another in a compound Usually a metal is exchanged for another metal or one nonmetal is exchanged for another nonmetal in an ionic compound Called a Neutralization reaction! Types of Chemical Reactions (still a double replacement) When acid and bases react they produce 5. Double Replacement Reactions water and a salt Exchanges occur between both reactants (which are both compounds!) Type of Reaction Reactants Products Formation (synthesis) A+ B → AB Simple Decomposition AB → A +B Single Replacement (A as a metal) A +BC → B + AC (D is a nonmetal) D + BC → C + BD Double Replacement AB + CD → AD +CB Neutralization Acid +Base → Salt + H2O Hydrocarbon Combustion C xH y + O 2 → CO2 + H2O Word Equations - #1 Hydrogen and oxygen react to produce water REACTANTS PRODUCTS 1. Write your skeleton equation! H2 (g) and O2 (g) → H2O (l) 2. Then BALANCE! 2 H2 (g) and O2 (g) → 2 H2O (l) Word Equations- #2 Remember what decomposition means? The decomposition of copper (II) oxide into copper and oxygen REACTANT PRODUCTS 1. Write your skeleton equation! CuO (S) → Cu (s) + O2 (g) 2. Then BALANCE! 2 CuO (S) → 2 Cu (s) + O2 (g) Word Equations - #3 Copper and silver nitrate react to produce copper (II) nitrate and silver REACTANTS PRODUCTS 1. Write your skeleton equation! Cu(s) + AgNO3 (s) → Cu(NO3)2 (aq)+ Ag (s) 2. Then BALANCE! Cu(s) + 2 AgNO3 (s) → Cu(NO3)2 (aq)+ 2 Ag (s) Word Equations - #4 Calcium nitrate and sodium hydroxide yield calcium hydroxide and sodium nitrate REACTANTS PRODUCTS 1. Write your skeleton equation! Ca(NO3)2 (s) +NaOH(s) → Ca(OH)2(aq) + NaNO3 (aq) 2. Then BALANCE! Ca(NO3)2 (s) + 2 NaOH(s) → Ca(OH)2(aq) + 2 NaNO3 (aq) Predicting Products Formation reactions ○ There will only be a single product ○ It will be an ionic compound (solid state) unless otherwise indicated ○ It must be written to create a neutral compound (follow the rules you learned previously for writing ionic compounds) ○ If the reactants and products are not equal, you must balance the equation, using coeﬃcients Practice Fe (s) + O2 (g) → Nitrogen and hydrogen react to produce ammonia Iron metal reacts with ﬂuorine gas to produce iron (III) ﬂuoride. Predicting Products Decomposition reactions ○ There will be two or more products ○ They will be elements ○ Elements are written alone using the state indicated on the periodic table, except for the diatomic and polyatomic elements (H2, N2, O2, F2, Cl2, Br2, I2, P4, S8) Practice NaO (s) → Na (s) + O2 (g) silver chloride decomposes lithium sulﬁde decomposes Single and Double Replacement Equations and Predicting Products Recap and Review → Single Replacement Reaction In a single replacement reaction, a reactive element reacts with an ionic compound. The element switches with an element in the compound so that you get an new compound and a new element element + compound → new compound + new element A + BC → B + AC Predicting Products Single Replacement Reactions ○ There will be two products ○ One will be a compound and the other an element ○ Elements are written alone using the state indicated on the periodic table, except for the diatomic and polyatomic elements (H2, N2, O2, F2, Cl2, Br2, I2, P4, S8) ○ Compounds are written using the ionic rules to make them neutral ○ The reactant element switches places with the element in the compound that forms the same type of ion (cation or anion) Practice - Single Replacement/ Predicting Ag (s) + H2S (g) → Cl2 (g) + NaBr (aq) → Potassium metal reacts with water Recap and Review - Double Replacement Reactions Double replacement reactions commonly occur between two ionic compounds in solution since ionic compounds are solid at room temperature This type of reaction often AB produces + CD →a precipitate AD + CB This is called a double replacement reaction because two new ionic compounds are formed Predicting Products AB + CD → AD + CB A and C are both cations- Will never pair up because ionic compounds consist of one positive ion and one negative ion A and C will always appear ﬁrst in formula because cations are always written ﬁrst B and D are anions, so they will combine with any positive ions-always written second in formulas check the solubility of both products Practice- Double Replacement/ Predicting NaCl (aq) + AgF (aq) → FeS (aq) + 2HCl (aq) → Barium chloride reacts with sodium sulfate Double Replacement/ Combustion Reactions Recap and Review - Double Replacement Reactions Double replacement reactions commonly occur between two ionic compounds in solution since ionic compounds are solid at room temperature This type of reaction often AB produces + CD →a precipitate AD + CB This is called a double replacement reaction because two new ionic compounds are formed Predicting Products AB + CD → AD + CB A and C are both cations- Will never pair up because ionic compounds consist of one positive ion and one negative ion A and C will always appear ﬁrst in formula because cations are always written ﬁrst B and D are anions, so they will combine with any positive ions-always written second in formulas check the solubility of both products Practice- Double Replacement/ Predicting NaCl (aq) + AgF (aq) → FeS (aq) + 2HCl (aq) → Barium chloride reacts with sodium sulfate Neutralization Reaction-Review Acid + Base → Salt + Water HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l) Acid reacts with base, to produce a salt and water They neutralize each other A neutralization reaction is also a double replacement reaction Practice- Neutralization Reaction LiOH (aq) + H2SO4 (aq) Hydroﬂuoric acid reacts with barium hydroxide Combustion Reaction - Review General formula for a hydrocarbon is CxHy –the x and y subscripts are whole numbers that indicate how many carbon and hydrogen atoms are in the molecule CxHy + O2 (g) → CO2 (g) + H2O(g) Ex: CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O(g) Predicting Products- Combustion Hydrocarbons are substances that contain hydrogen and carbon Ex: CH4(g), C3H8(g), C6H6(l) Combustion reactions always involve addition to oxygen These reactions are exothermic - they release energy If carbon is a reactant, CO2 (g) is a product If hydrogen is a reactant H2O (g) is a product Practice - Combustion Reactions C2H6 (g) + O2 (g) → Methanol Burns Neutralization Reactions/ Reactions Lab/ Start Moles Practice - Combustion Reactions C2H6 (g) + O2 (g) → Methanol Burns Neutralization Reaction-Review Acid + Base → Salt + Water HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l) Acid reacts with base, to produce a salt and water They neutralize each other A neutralization reaction is also a double replacement reaction Practice- Neutralization Reaction LiOH (aq) + H2SO4 (aq) Hydroﬂuoric acid reacts with barium hydroxide Reaction Lab- Videos/ Assignment Review your notes on how we know a chemical reaction has occurred! In this lab you will be watching videos of various chemical reactions Watch the associated video and then complete the questions that goes along with the video. For each chemical reaction: 1. STATE THE TYPE OF REACTION 2. BALANCE THE EQUATION 3. FILL IN THE MISSING STATES OF MATTER 4. RECORD YOUR OBSERVATIONS 5. ANSWER THE QUESTION THAT FOLLOWS Reaction Lab - Videos Q1: The burning of magnesium requires welders goggles to actually observe - Watch video: “Burning Magnesium Synthesis Video” - Answer questions Q2: Water molecules can be broken down by passing an electric current through an acidified water solution using a Hoffman Apparatus - Water video: “Electrolysis of water in a Hoffman Apparatus” - Answer questions Reaction Lab - Videos Q3: When you light a candle the nylon cord burns, but the reason why the candle burns for a long time is because the paraffin (candle wax) is also burning - Watch video: Lighting a candle reaction video - Answer questions Q4: Some aluminum foil has been left in a solution of copper (II) chloride - Water video: copper chloride and aluminum foil reaction video - Answer questions Reaction Lab - Videos Q5: The two compounds below are in a solution and mixed. What happens when they combine? - Watch video: calcium chloride and sodium carbonate reaction video - Answer questions Q6: Acetic acid (the main component of vinegar) and potassium hydroxide react…. - Water videos: - Acetic acid and potassium hydroxide reaction video - AP chemistry lab 11 - Bromothymol Blue as pH indicator video - Answer questions Reaction Lab - Videos Q7: When sodium is added to water, the sodium melts to form a ball and moves around the surface. It fizzes rapidly - Watch video: sodium and water reaction - Answer questions Please complete the lab on google slides, it has text boxes for you to type in. This lab is for MARKS and is DUE WEDNESDAY AT THE BEGINNING OF CLASS Moles and Molar Mass! What is a mole? A MOLE is a measure of QUANTITY Example: 1. ONE DOZEN = 12 2. PAIR = 2 3. ONE MOLE = 6.02 x 10 23 We can use it to count MOLECULES or ATOMS Where did we get the number 6.02 x 10 23 ? The mole is defined as the amount of a substance that contains as many elementary entities as exactly 12 g of carbon! ➔ This happened to be 6.02 x 1023 so that is the number that we use! ➔ Called AVOGADRO’S NUMBER Molar Mass of an Element Use your periodic table and locate the atomic molar mass (M) ○ UNITS: g/mol (or how many grams are in one mole) What is the molar mass of these elements? Molar Mass of a Compound To calculate the molar mass of a compound add the molar mass of each element in the ratio it is formed in the formula: Example: Calculate the Molar Mass of CO 2 M= 12.0l g/mol (CARBON) M= 2 (16.00 g/mol) (OXYGEN) M CO2 = 44.01 g/mol Practice! What is the molar mass of the following compounds? 1. H2SO4 2. Mg(NO3)2 Moles and Molar Mass! What is a mole? A MOLE is a measure of QUANTITY Example: 1. ONE DOZEN = 12 2. PAIR = 2 3. ONE MOLE = 6.02 x 10 23 We can use it to count MOLECULES or ATOMS Where did we get the number 6.02 x 10 23 ? The mole is defined as the amount of a substance that contains as many elementary entities as exactly 12 g of carbon! ➔ This happened to be 6.02 x 1023 so that is the number that we use! ➔ Called AVOGADRO’S NUMBER Molar Mass of an Element Use your periodic table and locate the atomic molar mass (M) ○ UNITS: g/mol (or how many grams are in one mole) What is the molar mass of these elements? Molar Mass of a Compound To calculate the molar mass of a compound add the molar mass of each element in the ratio it is formed in the formula: Example: Calculate the Molar Mass of CO 2 M= 12.0l g/mol (CARBON) M= 2 (16.00 g/mol) (OXYGEN) M CO2 = 44.01 g/mol Practice! What is the molar mass of the following compounds? 1. H2SO4 2. Mg(NO3)2 Mole Formula m = mass in grams(g) n = # of mols n= m M M = Molar Mass of compound (found on periodic table) Practice! 1. If the mass of an iron bar is 16.8 g, what is the molar amount n= ? m= 16.8g n= m M (iron) = 55.85 g/mol M = 16.8 g 55. 85 g/mol = 0.3008 mol = 0.301 mol Practice! 1. What is the molar amount of 10g of CO2? n= ? m= 10 g n= m M (iron) = 44.01 g/mol M M of CO2 12.01 g/mol + 2(16.00 g/mol) = 44.01 g/mol Manipulating the Formula! - NO TRIANGLE!!!! Note: To manipulate the formula to solve for m: n= m m = nM M Example: What is the mass of 7.50 mol of H2O? m= ? n= 7.50mol M= (H2O) = (2 (1.01) + (16.00) = 18.02 g/mol m= nM CHEMISTRY DONE!!!

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