Chemistry Notes - Oxidation Numbers & Naming Molecules PDF

Oxidation numbers What are Oxidation Numbers?  Oxidation numbers represent the charge an atom would have if electrons were transferred completely in bonds.  Oxidation numbers help us track electron movement during chemical reactions, especially in redox reactions. Key Rules for...

Oxidation numbers What are Oxidation Numbers?  Oxidation numbers represent the charge an atom would have if electrons were transferred completely in bonds.  Oxidation numbers help us track electron movement during chemical reactions, especially in redox reactions. Key Rules for Assigning Oxidation Numbers  1. The oxidation number of an element in its pure form (e.g., O₂, N₂, H₂) is 0.  2. For monoatomic ions, the oxidation number equals the ion's charge (e.g., Na⁺ = +1, Cl⁻ = -1).  3. Oxygen has an oxidation number of -2, except in peroxides (-1).  4. Hydrogen is +1 when bonded to non-metals and -1 when bonded to metals.  5. The sum of oxidation numbers in a neutral compound is 0, and in polyatomic ions, it equals the ion's charge. Example: Assigning Oxidation Numbers (H₂SO₄)  In H₂SO₄ (Sulfuric Acid):  Hydrogen (H) = +1 (Rule for hydrogen)  Oxygen (O) = -2 (Rule for oxygen)  Let sulfur (S) be x.  The sum of oxidation numbers must be 0:  2(+1) + x + 4(-2) = 0  2 + x - 8 = 0 → x = +6  Oxidation number of sulfur (S) = +6 What is Oxidation and Reduction?  **Oxidation**: Loss of electrons or an increase in oxidation number.  **Reduction**: Gain of electrons or a decrease in oxidation number.  The element undergoing oxidation is the reducing agent and the element undergoing reduction is the oxidizing agent.  In a redox reaction, one species is oxidized (loses electrons) while another is reduced (gains electrons). Example: Identifying Oxidation and Reduction (Zn + CuSO₄)  In the reaction: Zn + CuSO₄ → ZnSO₄ + Cu  Zn is oxidized: Zn → Zn²⁺ + 2e⁻ (Oxidation, because Zn loses electrons)  Cu²⁺ is reduced: Cu²⁺ + 2e⁻ → Cu (Reduction, because Cu²⁺ gains electrons)  Zn undergoes oxidation and Cu²⁺ undergoes reduction. Tips for Identifying Oxidation and Reduction  If the oxidation number of an element increases, it is oxidized.  If the oxidation number decreases, it is reduced.  Use the rules for assigning oxidation numbers to determine changes in a reaction.  Redox reactions always involve both oxidation and reduction happening together. Extra Practice Assign the oxidation number for the elements below and identify which element is oxidized and which element is reduced. 1- : 2H2+O2→ H2O 2- Zn+2HCl→ZnCl2+H2 3- Fe2O3+3CO→2Fe+3CO2 4- 2Na+Cl2→2NaCl Naming molecules Naming ionic compound As a revision to name ionic compound please watch the following video using the link below https://www.youtube.com/watch? v=CkCzceecCrc Note that you need to memorize the following list to be able to name compounds that contains polyatomic ions Naming binary and molecular compouds  Binarycompounds : Binary compounds are compounds that contains two non- metals example : CO , CO2 , NH3 How do we name Binary Compounds  Naming binary molecular compounds - Binary molecular compounds are composed of molecules containing two non- metal that share electrons - The names of the binary molecular compounds give both the number and the type of each atom present. - Prefix element name Prefix element name with ‘ide’ ending - Greek prefixes are used to indicate the number of each atom present in a molecule of the compound. Naming Binary Molecular Compounds Examples Name the following binary compounds  A- SO2  B. NF3  C. CCL4 A. SO2 : Sulfur dioxide it is not mono sulfur dioxide we do not use the prefix mono with the first element. B. NF3 Nitogen Triflouride C. CCL4 :Carbon tetrachloride  For extra practice on naming Binary compounds please watch The video below  https://www.youtube.com/watch?v=59xC6c2TV4w Binary Acids - Binary acids are acids that contains Hydrogen and another non metal example :  HF , HBr , HCl , HI. Naming Binary acids  Binary acids - These are acids that have only 2 different types of atoms in the formula. - Hydrogen is always one of the atoms - The other atom is always a simple non-metal. - The binary acid name always starts with the prefix hydro – and ends with the name –ic acid. Naming binary acids  HI : hydroiodic acid  HF : hydroflouric acid  H2S. : hydrosulfuric acid. Oxyacids  Oxyacids - Must contain hydrogen - Must have a polyatomic ion with oxygen in it - Examples - HCLO3 , HCLO4 , H2SO4 Naming Oxyacids 1. Identify the polyatomic ion ( anion) the polyatomic ion in this case will always contain oxygen. 2. if the polyatomic ion name ends in – ate- then change ending to – ic suffix. 3. if the polyatomic ion name ends in - ite - the change ending to ous suffix 4. write word acid at the end of all names. Name the following acids - H2SO4 This is an oxyacid so the name depends on the oxyanion The oxyanion in the acid above is sulfate so its name ends in ate replace ate with –ic and add the word acid so H2SO4 IS SULFURIC ACID - HCLO2 This is an oxyacid the oxyanion is called chlorite it ends with –ite so replace – ite with –ous and add the word acid the name of the acid above is chorus acid Naming oxyacids  You can refer to the video below to practice more how to name binary acids and oxyacids. https://www.youtube.com/watch?v=z6oylZ7KBX0 CHAPTER 9 LESSON 1 Representing chemical reactions Representing chemical reactions  Chemists use statements called equations to represent chemical reactions  Reactants are the starting substances  Products are the substances formed in the reaction  This table summarizes the symbols used in chemical equations  Word Equations  A word equation describes chemical change using the names of the reactants and products.  Write the word equation for the reaction of methane gas with oxygen gas to form carbon dioxide and water.  Skeleton Equations  A skeleton uses chemical formulas rather than words to identify the reactants and the products.  Ex iron(s). + Chlorine (g) → iron (III) chloride(s)  Fe(s) + Cl(g) → FeCl3 (s)  C(s) + S(s) → CS2 Skeleton Equations  Give you the chemical formula of the reactants and products  Al(s) +. Br 2(I) + AlBr 2(s)  The small letter in parenthesis tells you what state the elements of the compound are in.  Skeletal equations give the same amount of information as word equations, but the skeletal equations show you more at a glance. Balancing chemical equations  Pleasewatch this video that explains the steps for balancing chemical equations. https://www.google.com/search?q=Balancing+chemical+equ ations+ppt&source=lmns&safe=strict&hl=en&ved=2ahUKEwi hncmY1p_oAhU Why do we need to balance a chemical equation? - Sometimes , a written equation Does not have equal number of reactants and products. - For every chemical reaction ,the law of conservation of Mass must be followed. - The law pf conservation of mass states : Atoms can be created or destroyed. - Therefore ,the number of atoms in the reactants must equal the number of protons. Balancing chemical Equations Roles of balancing Chemical Equations Please solve the practice problems page 287 Classifying chemical reactions Writing formulas for ionic compounds  Write the formula for the cation and anion found in the compound taking into consideration the charge of each ion found.  If the charges are equal to each other they cancel each other and we can write directly the formula of the compound  Example Magnesium oxide Writing the formulas of ionic compounds  if the charges of the ionic compounds are different from each other , crisscross the charges to identify the combining ratios.  Always write the cation first and then the anion and never show the charges as a part of the chemical formula since the unit being represented is electrically neutral. Writing the formulas of ionic compounds Note that you need to memorize the following list to be able to write the formulas of the compounds that contains polyatomic ions What is chemical reaction?  The change of one or more substances into other substances having different compositions and properties is called a chemical reaction.  Example  C(s) + O2 (g) → CO2 (g)  2H2 + O2 (g) → 2H2O (g)  In a chemical reaction the substance which react together are called reactants whereas the new substances forms are called products.  Reactants → Product Types of chemical reaction There are five main types of chemical reactions that we will be covering in this chapter : - Combination reactions - Decomposition reactions - Single –replacement reactions - Double –replacement reactions - Combustion reactions  Synthesis .Two or more substances react to produce a single new substance  A + B → 𝐴𝐵 - Synthesis of two elements C + O2 → 𝐶𝑂2 - Synthesis of two compounds 2CaO + 2H2O → 2𝐶𝑎 𝑂𝐻 2 - Synthesis of a compound and an element 2C0 + O2 → 2𝐶𝑂2 Decomposition reaction  Decomposition reactions occur when a compound breaks up into elements or in a few to simpler compounds.  Combustion Reactions  Definition Reaction where an element or a compound reacts with oxygen ,often producing energy in the form of heat and light,  Examples  CH4 + 2O2 → CO2 + 2H2O + heat + light  2Mg(s) + O2(g) → 2MgO(s)  Note that the second combustion reaction can also be classified as synthesis reaction. However not all combustion reactions are synthesis reactions Practice problems page 291 Practice problems page 292 Replacement reactions  A simple –replacement reaction is a chemical change in which one element replaces a second element in a compound.  All single replacement reactions have the genera form/  A +. BC → B + AC  For example  Zinc metal + aqueous copper nitrate.  The predicted products are the element Cu and the compound Zn(NO3)2  Reaction equation Zn (s) + Cu(NO3)2(aq) → Cu(s) + Zn(NO3)2 (aq) 1- A metal replaces hydrogen or another metal : Example Cu + AgNO3 → ?? To predict the products first we should check the free metal present if it is a metal it will replace another metal or a hydrogen atom. In the example above for cu to replace Ag it should be more reactive than it. To know which metal is more reactive than the other check the reactivity series below: For the above example and depending on the reactivity series below we can tell that Cu is more reactive than Ag so Cu can replace Ag. The reaction will be given by : Cu + 2AgNO3→ 2Ag + Cu(NO3)2 However if the following reaction was given Ag + Cu(NO3 ) → ?? The free element is Ag which is metal so it will replace another metal so it will replace another metal which is Cu. Ag is less reactive than Cu So no reaction will take. Ag + Cu(NO3 ) → No reaction  Nonmetal replaces nonmetal.A third type of single replacement reaction involves the replacement of a nonmetal in a compound by another nonmetal. Halogens are frequently involved in these types of reactions. like metals ,halogens exhibit different activity levels in single –replacements reactions. The reactivities of halogens determined by single replacement reactions are also shown in figure 13.The most active halogen is fluorine and the least active is iodine.A more reactive halogens replaces a less reactive halogen that is part of a compound dissolved in water.  For example fluorine replaces bromine in water containing dissolved sodium bromide. However bromine does not replace fluorine in water containing dissolved sodium fluoride.  F2(g) + 2 NaBr(aq) → 2NaF(aq) +Br (l)  Br2(g) + 2NaF(aq) → NR Practice problems page 295 Section Review page 298  Chapter 6 , section 2 Classification of the elements Organizing the elements by Electron configuration: Remember that electron configuration determines the chemical properties of an element. An element’s electron configuration and its number of valence electrons can be determined from its position in the periodic table. Organizing the elements by Electron configuration: Electron configuration : Recall that the electron configuration is the arrangement of electrons in an atom. it tends to assume the arrangement that gives the atom the lowest possible energy. It is the most stable arrangement and is called the ground state configuration. 3 rules ( The aufbau principle , the pauli exclusion principle and Hund’s rule ) define how electrons can be arranged in atom’s orbital. Organizing the elements by Electron configuration: Electron configuration : Organizing the elements by Electron configuration Valence electrons : recall that electrons in the highest principle energy level of an atom are called valence electrons. Elements in the same group have similar chemical properties because those elements have the same number of valence electrons. Organizing the elements by Electron configuration One can determine an atom’s electron configuration and its number of valence electrons from its position on the periodic table. Organizing the elements by Electron configuration Valence electrons and the period Energy level of an element’s valence electrons indicates the period on the periodic table in which it is found For example Lithium=period 2 since the valence electrons were in the second energy level Organizing the elements by Electron configuration Organizing the elements by Electron configuration The s- p- d- and F-blocks elements The periodic table is divided into sections or blocks representing the atom’s energy sublevel being filled with valence electrons There are 4 energy sublevels. Organizing the elements by Electron configuration Organizing the elements by Electron configuration P-block elements Compromised of group 13-18 P-block span 6 groups because it can hold 6 electrons Group 18=unique member of the p-block They are so stable=rarely ever had chemical reactions This is because both their s- orbital and their p-orbitals are completely filled S- and p- blocks make up representative elements. Organizing the elements by Electron configuration Organizing the elements by Electron configuration d- block elements Contains transition metals Usually characterized by - Filled outermost s orbitals on energy level n - Filled or partially filled d orbitals of energy level n-1 - As you move across the periods the d- orbitals are filled Organizing the elements by Electron configuration F- block elements Contain the inner transition metals Characteristics : Filled or partially filled outermost s orbital Filled or partially filled 4f and 5f orbitals Electrons do not fill the f-orbitals in any predictable manner. F- orbitals can hold 14 electrons=14 groups on the periodic table. Organizing the elements by Electron configuration Organizing the elements by Electron configuration Organizing the elements by Electron configuration Chapter 6 Lesson 3 The periodic law : Atoms are arranged by increasing atomic number. The elements display a repeating pattern of chemical and physical properties. Periodic Trends : Periodic trends are patterns that are present in the periodic table. They show different properties of elements , and how these characteristics increase or decrease as you move across a row or down a column of the periodic table. Four main periodic trends : 1- Atomic Radius 2- Ionization Energy 3- Electronegativity 4- Ionic radius What is the trend for atomic radius as we go down a group? Vertical trend : As we go down a family , valence electrons are found in energy levels further from the nucleus. - As we go down a column the atomic radius will increase Quick practice Which one has a larger atomic radius ? O or S Ca or Be What is the horizontal trend for atomic radius? ( Trend across a period) As we move across a row ,we are still in the same energy level however the atomic number is increasing causing an increase in the effective nuclear charge.so the nucleus is stronger and it can attract more the outer most energy level causing a decrease in the atomic radius. Quick Practice Which atom is smaller? Mg or Si Ga or As 2- Ionization Energy What is ionization energy? Ionization energy is the energy required to remove an electron from an atom. It is measured in kilojoules KJ Ionization energy and atomic radius are inversely proportional. The larger the atom is the easier its electrons are removed. The amount of energy required to remove an electron depends on shielding. Shielding is when the core (inner) electrons block the protons from pulling on the valence electrons Ionization energy across column: As we go down a group the atomic radius increases so the amount of ionization energy needed to remove the electrons decreases. Ionization energy across a period As we go across a row , the atomic radius becomes smaller which means that the ionization energy will increase. -  The second ionization energy I.E is the energy required to remove the second electron(s). Always greater that first I.E. The third I.E is the energy required to remove a third electron, which is greater than 1st and 2nd I.E. Successive values of Ionization Energy: If you examine the previous table , you will notice that successive ionization energy for a certain element increases. From the previous table you will notice that the ionization energy at which the large increase in energy occurs is related to the number of valence electrons. For example magnesium has two valence electrons because the dramatic increase in ionization energy occurs after the second I.E value. 1. Element X has two valence electrons because between second and third ionization energy there is a dramatic increase in the IE. 2. Element Y has 4 valence electrons because after the third IE there is a dramatic increase in the IE values. 3. Element Z has three IE because after the third IE there is a dramatic increase in the IE values. 3- Electronegativity : What is electronegativity ? Electronegativity of an element indicates the relative ability of its atoms to attract electrons in a chemical bond. Electronegativity down a group : As we move down a group the electronegativity decreases. The electronegativity down a group decreases because the atomic radius increases so the ability of an atom to gain electrons decreases. Electronegativity across period : As we move from left to right in the periodic table , the electronegativity increases. Unit for electronegativity Melting and boiling point : Melting point : The melting point is the amount of energy needed to break a bond in order to change a substance from a solid state to a liquid state. Boiling point : The boiling point is the amount of energy required to break a bond to change a substance from a liquid state to the gas state. The size of the radius of the ion depends on the overall charge of the ion positive or negative 1- anions are larger than their respective atoms. Electron-electron repulsion forces them to spread apart Electrons outnumber protons The protons cannot pull the extra electrons more tightly toward the nucleus. 2- cations are smaller than their respective atoms Less electron-electron repulsion allows for them to come closer together Protons outnumber electrons The protons can pull the fewer electrons toward the nucleus more tightly If am electron that is lost is the only valence electron causing the electron configuration of the cation to be like a noble gas than the entire energy level is lost in this case the radius of the cation is much smaller than its respective atom.  CHEMISTRY NOTES The reactivity series In a reactivity series, the most reactive element is placed at the top and the least reactive element at the bottom. More reactive metals have a greater tendency to lose electrons and form positive ions. A reactivity series of metals could include any elements. For example: A good way to remember the order of a reactivity series of metals is to use the first letter of each one to make up a silly sentence. For example: People Say Little Children Make A Zebra Ill Constantly Sniffing Giraffes. Observations of the way that these elements react with water, acids and steam enable us to put them into this series. The tables show how the elements react with water and dilute acids: ] Element Reaction with water Potassium Violently Sodium Very quickly Lithium Quickly Calcium More slowly Element Reaction with dilute acids Calcium Very quickly Magnesium Quickly Zinc More slowly Iron More slowly than zinc Copper Very slowly Silver Barely reacts Gold Does not react Note that aluminum can be difficult to place in the correct position in the reactivity series during these experiments. This is because its protective aluminum oxide layer makes it appear to be less reactive than it really is. When this layer is removed, the observations are more reliable. Displacement reactions of metal oxides A more reactive metal will displace a less reactive metal from a compound. The thermite reaction is a good example of this. It is used to produce white hot molten (liquid) iron in remote locations for welding. A lot of heat is needed to start the reaction, but then it releases an incredible amount of heat, enough to melt the iron. aluminum + iron(III) oxide → iron + aluminum oxide 2Al + Fe2O3 → 2Fe + Al2O3 1  ] Because aluminum is more reactive than iron, it displaces iron from iron(III) oxide. The aluminum removes oxygen from the iron (III) oxide: iron is reduced aluminum is oxidized Reactions between metals and metal oxides allow us to put a selection of metals into a reactivity series. Using metals A, B and C: Metal B Metal C Metal A A oxide X Displaces A Displaces A B oxide No reaction X No reaction C oxide No reaction Displaces C X Metal A cannot displace either B or C - so it must be the least reactive and be at the bottom of this reactivity series. Metal B displaces both A and C - so it must be the most reactive and be at the top of this reactivity series. Metal C displaces A but cannot displace B - so it must be more reactive than A but less reactive than B, and be in between them in this reactivity series. In general, the greater the difference in reactivity between two metals in a displacement reaction, the greater the amount of energy released. Aluminum is much higher than iron in the reactivity series, so the thermite reaction releases a lot of energy. Magnesium is very high in the reactivity series, and copper is very low - so the Reaction between magnesium and copper oxide is more violent. Therefore, the order is: 2 ] Displacement reactions of solutions A more reactive metal will displace a less reactive metal from a solution of one of its salts. For example: Magnesium + copper (II) sulfate → copper + magnesium sulfate Mg(s) + CuSO4 (aq) → Cu(s) + MgSO4(aq) In this reaction, the blue color of the copper (II) sulfate fades as it is used up (magnesium sulfate solution is colorless). We would also see copper metal forming. Reactions between metals and solutions of metal salts allow us to put a selection of metals into a reactivity series. Using metals J, K and L: Metal K Metal L Metal J J sulfate X No reaction No reaction K sulfate Displaces K X Displaces K L sulfate Displaces L No reaction X Metal J displaces both K and L - so it must be the most reactive and be at the top of this reactivity series. Metal K cannot displace either J or L - so it must be the least reactive and be at the bottom of this reactivity series. Metal L displaces K but cannot displace J - so it must be more reactive than K but less reactive than J, and be in between them in this reactivity series. Therefore, the order is: 3  ] Oxidation and reduction Oxidation is the loss of electrons from a substance. It is also the gain of oxygen by a substance. For example, magnesium is oxidized when it reacts with oxygen to form magnesium oxide: Magnesium + oxygen → magnesium oxide 2Mg + O2 → 2MgO Reduction is the gain of electrons by a substance. It is also the loss of oxygen from a substance. For example, copper (II) oxide can be reduced to form copper when it reacts with hydrogen: Copper (II) oxide + hydrogen → copper + water CuO + H2 → Cu + H2O Usually, oxidation and reduction take place at the same time in a reaction. We call this type of reaction a redox reaction. Note that: the oxidizing agent is the chemical that causes oxidation the reducing agent causes the other chemical to be reduced Take a look at the following thermite reaction: Aluminum + iron (III) oxide → iron + aluminum oxide It is easy to see that the aluminum has been oxidized. This means that the iron oxide is the oxidizing agent. We can also see that the iron oxide has been reduced. This means that the aluminum is the reducing agent. 1. 4  ] Rusting Rusting is an oxidation reaction. The iron reacts with water and oxygen to form hydrated iron (III) oxide, which we see as rust. Here is the word equation for the reaction: Iron + water + oxygen → hydrated iron (III) oxide Iron and steel rust when they come into contact with water and oxygen. Both water and oxygen are needed for rusting to occur. In the experiment below, the nail does not rust when air (containing oxygen) or water is not present: Aluminum does not rust (corrode) because its surface is protected by a natural layer of aluminum oxide. This prevents the metal below from coming into contact with air (containing oxygen). Unlike rust, which can flake off the surface of iron and steel objects, the layer of aluminum oxide does not flake off. Preventing rusting There are several ways to prevent iron and steel rusting. Some of these work because they stop oxygen or water reaching the surface of the metal: oiling - for example, bicycle chains greasing - for example, nut and bolts painting - for example, car body panels coating with a thin layer of plastic Iron and steel objects may also be covered with a layer of metal. Food cans, for example, are plated with a thin layer of tin. 5 ] Galvanizing Galvanizing is a method of rust prevention. The iron or steel object is coated in a thin layer of zinc. This stops oxygen and water reaching the metal underneath - but the zinc also acts as a sacrificial metal. Zinc is more real sacrificial protection A reactivity series lists metals in order of how reactive they are. Magnesium and zinc are often used as sacrificial metals. They are more reactive than iron and lose their electrons in preference to iron. This prevents iron from losing its electrons and becoming oxidized. 1. 6