Gr 10 Chemistry Unit PDF

Gr 10 Chemistry Unit Ionic compounds Characteristics: ​ Composed of cation (+) and anion (-) ​ Compound is held together by electric attraction (aka positive and negative charges) ​ Arranged in an alternating pattern to form an ionic crystal ​ Soluble in water (since it is polar), ion...

Gr 10 Chemistry Unit Ionic compounds Characteristics: ​ Composed of cation (+) and anion (-) ​ Compound is held together by electric attraction (aka positive and negative charges) ​ Arranged in an alternating pattern to form an ionic crystal ​ Soluble in water (since it is polar), ions dissociate ​ Ionic compound dissolves in a solvent because ions separate ​ Insulator in solid form, conductor once dissolved in water ○​ Conductivity = how easily it is for electricity to flow ○​ Water molecules pull ionic compounds apart (they are dissociated) ○​ They are now dissolved in water as positive and negative ions ​ Strong electrical attraction, difficult to melt the compound ○​ High boiling & melting point Lewis Diagram: ​ Draw valence electrons around the symbol ○​ For cations that lose electrons to become full, no electrons have to be drawn ○​ For anions that gained electrons, draw original electrons with dots and new ones with “x” ○​ Add bracket to show that it is an ion ○​ Write the charge on the top right corner (eg. 2+ ) ○​ ○​ Always write metal than non metal Three types of ionic compounds: 1.​ Binary Ionic Compounds (single charge metal + non metal) 2.​ Ionic Compounds with multivalent metals (transitional metals + non metal) 3.​ Ionic Compounds with Polyatomic metals Zero Sum Rule: ​ “The overall charge of an ionic compound must equal to zero” Binary Ionic Compounds: ​ Naming ○​ Name of metal + name of non metal w/ -ide ending ○​ Eg. Magnesium fluoride ​ Criss-cross rule ○​ Shortcut to determine the chemical formula of ionic compounds ○​ 1. Write ion charges above symbol ○​ 2. Criss-cross the charge numbers so they become subscripts ​ Note: ​ Subscripts must be in simplest form (reduced) ​ Metal comes first in the formula ​ Multivalent Compounds: Multivalent Compounds - Ionic compounds that involve transitional metals that can become a stable ion in more than one way. ​ Must specify which ion is used in the compound when naming these elements. ​ Refer to the table for which ion ​ Ions can have very different characteristics  ​ Chemical formula: ○​ Name tells which ion to use ○​ Eg. tin (IV) = sn with 4+ charge ​ Naming: ○​ Name of metal + ion charge in bracket + name of non metal w/ - ide ending ○​ 1. Find the non metal with the already known charge ○​ 2. Reverse criss cross ​ Note: ​ Subscripts might have been reduced to simplest form ​ Have to “unreduce” it Polyatomic Ions ​ Polyatomic Ions - An ion made up of more than one atom that acts as a single particle ​ Refer to the chart for list of polyatomic ions ​ Naming: ○​ Do not change into -ide endings, refer to chart ​ Chemical formula: ○​ Polyatomic ions are a single symbol ○​ Add second subscript outside of bracket ○​ Only add brackets if there is a second subscript Covalent Compounds Characteristics: ​ Formed between 2 nonmetals ​ Atoms share electrons instead of transfer ​ No ions are formed, the molecules is held together by covalent bonds ​ Covalent bond between atoms are very strong ​ Molecule = a group of atoms held together by covalent bonds ​ Forces of attraction between molecules weak ○​ Therefore low melting point ​ Cannot conduct electricity because they don’t have ions Naming: ​ Element 1 + element 2 -ide ending + prefix between each element (to indicate amount of atoms in the compound) ○​ Prefixes: Mono - 1 Di - 2 Tri - 3 Tetra - 4 Penta - 5 Hexa - 6 Hepta - 7 Octa -8 ○​ Drop “o” or “a“ in the prefix if used in front of element that begins with vowel ○​ Do not use “mono” prefix in front of the first element ○​ For first element, no prefix = assume 1 Chemical formula: ​ Use prefix to determine number of atoms ​ Don’t crisscross nor reduce ○​ Crisscross is only for finding out how many atoms are needed (subscripts) ○​ Eg. ​ N and H ​ N = 3 - ​ H = 1 + ​ Therefore, Nitrogen Trihydride, NH 3 is formed Lewis Diagram: ​ Share bonds ​ Location of electrons are different from how it is in Ionic Compounds ​ Structural diagram: ​ Easier to tell what is being shared, what is not ​ ​ Line = two electrons that are being shared Diatomic molecules: ​ 7 elements where their atoms are found naturally as pairs ​ HOFFBrINCl ​ Stabilize themselves by sharing electrons  ​ ​ Held together by a covalent bond Comparing covalent compounds and ionic compounds Binary Ionic Compounds Covalent Compounds Formed Cation (metals) and anion (non 2 non metals between metals) Why do they do Transferred from metal to non metal Share electrons with their valence electrons How does the Positive/ negative attraction holds Covalent bond compound stay ion together What would the Crystal lattice of positive and One single molecule particle look negative ions like What happens Ions come apart Ions stay together when compound dissolve in water Chemical Reactions In a chemical reaction, the total mass of the reactants is equal to the mass of the products. Matter is not created or destroyed. - Conservation of mass Types of chemical reactions: Synthesis, decomposition, combustion, single replacement, double displacement. (Can be ionic or covalent) When predicting products, determine subscript by charge of the other element (the element one element is combining with) then balance after “Rule” for single displacement and decomposition: Metal Carbonate → CO2 + Metal Oxide Metal Hydroxide → H2O + Metal Oxide Synthesis Definition: ​ When 2 or more reactants combine to form one product ​ A + B = AB ​ Eg. ○​ Sodium + Chlorine → Sodium Chloride ​ Na + Cl → NaCl ○​ Haber Process ​ N2 + 3 H2 → 2 NH3 Decomposition Definition: ​ When compound breaks into its component parts ○​ Eg. Glucose → Ethanol + Carbon dioxide ​ C6 H12 O6 → 2 C2 H6 O + 2 CO2 ​ Fermentation - common decomposition reaction when substances are broken down by bacteria and yeast ○​ Eg. Alcohol, cheese, yogurt Single replacement Definition: ​ Occurs when one element is replaced by another in a compound ​ A + BC → AC + B ​ Occurs when A is more reactive than B, giving a more stable product ​ Refer to the reactivity chart for who “kicks” who out, the greater the distance between the element the more vigorous the reaction ○​ Some equations might be “not reactive”, happens when the single element is not as reactive as the one in the compound ​ Eg. Steel ○​ 3CO + Fe2 O3 → 2 Fe + 3 CO2 Double displacement Definition: ​ AB + CD → AD + CB ​ When parts of two ionic compounds are exchanged ​ Metals and nonmetals will only replace other metals and nonmetals ​ Acid and bases is a type of double replacement ​ Eg. ○​ Magnesium Hydroxide (antacid) and Hydrochloride (stomach acid) ​ Mg (OH)2 + 2 HCl → MgCl2 + 2 H2O Combustion Definition: ​ Where a molecule (usually organic), reacts with oxygen ​ Equation = organic compound + oxygen → carbon dioxide + water ​ Organic molecules - Contains carbon-hydrogen bonds ​ Eg. Propane, used in BBQ grills, stoves ○​ C3 H8 + 5O2 → 3 CO2 + 4 H2O Combustion reaction: ​ When fuel “burns” with oxygen ​ Reactant is usually hydrocarbon, made up of only H and C ​ Bunsen burner ○​ Closed air hole = limited O2 = incomplete combustion = yellow flame  ○​ Open air hole = unlimited O2 = complete combustion = blue flame ○​ Blue flame is more energy efficient bc more heat is produced with same amount of gas Complete combustion ​ Depending on the availability of oxygen ​ When oxygen is plentiful: ○​ Hydrocarbon + oxygen → carbon dioxide + water + energy ○​ Methane (CH4) is a common hydrocarbon ​ CH4 (g) + 2O2(g) → 2H2O (g) + CO2 energy Incomplete combustion ​ When oxygen supply is limited ​ Products are carbon monoxide gas (CO) and carbon (C) in addition to carbon dioxide and water ​ Carbon produced in incomplete combustion are soots ​ Carbon monoxide is a highly toxic gas ○​ Odourless, colourless ○​ Symptoms of carbon monoxide poisoning ​ Headache ​ Dizziness ​ Nausea ​ Orange flickering flame indicates incomplete combustion ​ C4H10 (g) + 5O2 (g) → 2CO2(g) + CO (g) + C (s) + energy Combustion of Hydrogen ​ 2H2 + O2 → 2H2O + energy ​ Benefit: ○​ Burns cleanly, produces only water and energy ○​ Endless water to produce hydrogen ​ Drawbacks: ○​ Producing hydrogen requires energy ○​ Hydrogen fueled energies are expensive ○​ Hydrogen is explosive Acid and Bases Bronsted-Lowry Theory: ​ Acids donate a proton (H+) ​ Bases accepts a proton (H+) ​ Water = amphoteric (both an acid and a base) Electrolytes: ​ Species that conduct electricity when dissolved ​ Acids, bases and salts are electrolytes ​ Salt and strong acids of bases form strong electrolytes ○​ They are fully dissociated therefor all ions are available to conduct electricity ​ Weak acids and weak bases from weak electrolytes ○​ They are partially dissociated, some are still in molecular form ○​ Covalent bonds are strong in molecular form, can’t be broken up to conduct electricity Base: ​ Solution that has an excess of OH- ions ​ Another word for base is alkali ​ Bases accepts hydrogen ions ​ Properties: ○​ Feel slippery ○​ Taste bitter ○​ Corrosive ○​ Can conduct electricity ○​ Do not react with metals ○​ Turns orange litmus paper purple ​ Uses of base: ○​ Give soaps, ammonia and cleaning products their useful properties ○​ OH- ions interact strongly with dirt and grease ○​ Chalk and oven cleaner are examples of products containing bases ○​ Blood Acid: ​ Solution that has an excess of H+ ions ​ The more H+ ions, the more acidic ​ Properties: ○​ Sour ○​ Conducts electricity  ○​ Corrosive ○​ Some reacts strongly with metals ○​ Turns orange litmus paper red ​ Use of acids: ○​ Acetic acid = vinegar ○​ Citric acid = lemons, limes ○​ Ascorbic acid = vitamin C ○​ Car batteries are made from acids ○​ Sulfuric acid is used to produce fertilizers, steel, paints, plastics Acid base reactions ​ Neutralization ○​ The acid-base mixture is not as acidic/basic as the individual starting solutions ○​ One acid and one base as reactants ○​ Produces water and ionic salt ○​ Acid (aq) + Base (aq) → water (l) + salt (aq) (*Notice how everything is aqueous except for water) ​ Neutralization with CO3 or HCO3 ○​ Breaks down into H2O (l) + CO2 (g) ​ Common salts formed by acid-base reactions ○​ Sodium chloride/ potassium chloride = salt in food ○​ Calcium chloride = de-icer for roads ○​ Calcium carbonate = found in limestone ○​ Ammonium nitrate NH4 NO3 = fertilizer ​ PH scale ○​ Measure how acidic or basic a solution is ○​ Ranges from 0-14 ○​ 7 = neutral ​ Pure water, blood plasma, milk ○​ Below 7 = acidic ​ Hydrochloric, vinegar, soft drinks ○​ Above 7 = basicl ​ Ammonia, soap, bleach, concentration of alkali solution ​ PH indicators ○​ Phenolphthalein (turns from colourless to pink if PH range is between 8.2-10) ○​ Baking soda (reacts with acid) ○​ Bromothymol blue (turns from yellow to blue if PH range is between 6.0-7.6)  ○​ Litmus paper (turns from orange to red if acidic, orange to purple if basic) Acid nomenclature ​ Formula of acids always start with Hydrogen ​ Acids always has a subscript (aq) since it is dissolved in water ​ Binary acids ○​ Hydrogen + nonmetal element ○​ Has a general formula H#X(aq) ○​ Naming: ​ Hydro + nonmetal element prefix + ic ending + acid ​ Eg. H2S(aq) = hydrosulfuric acid ​ Oxyacids ○​ Hydrogen + polyatomic containing oxygen ○​ Naming: ​ Polyatomic ion prefix + ic + acid ​ Eg. H2SO4 (aq) = Sulfuric acid ​ Giving formulas: ○​ Determine whether it is binary or oxyacid ​ Oxyacid = without “hydro” ○​ Crisscross hydrogen with element/polyatomic ion ​ Bases: ○​ Always contains hydroxide group ○​ Name and give formula as regular ionic compound