Chemical Formulae for Elements and Compounds PDF

Chemical formulae for elements The chemical formula of a number of elements is simply its’ chemical symbol e.g. Element name Chemical formula Magnesium Mg Iron...

Chemical formulae for elements The chemical formula of a number of elements is simply its’ chemical symbol e.g. Element name Chemical formula Magnesium Mg Iron Fe Phosphorus P Argon Ar However some elements exist in nature as two atoms joined together. These are known as the Diatomic elements. There are 7 diatomic elements. Whenever we write the chemical formula of these elements we put a subscript ‘2’ beside them. This means that each molecule has two atoms joined together. The seven diatomic elements are: Element name Chemical formula Iodine I2 Oxygen O2 Chlorine Cl2 Bromine Br2 Nitrogen N2 Hydrogen H2 Fluorine F2 Chemical Formulae for Compounds The chemical formulae for simple compounds containing two elements can be found using the valency of the composing elements. The names of these elements sometimes end in ‘ide’, for example NaCl, the compound between the elements sodium (Na) and Chlorine (Cl) is named as sodium chloride. Note that the name of some ions can also end in ‘ide’, for example, OH- is named Hydroxide and CN- is named cyanide. The valency depends on the group number as follows: Group 1 2 3 4 5 6 7 8 Number Valency 1 2 3 4 3 2 1 0 We can follow a set of simple rules to write chemical formula of compounds containing two elements: 1. Write symbols of elements present in compound 2. Put valency above each symbol 3. Cross valency over (swap and drop) 4. Cancel down ratio if necessary 5. Write correct chemical formula Example: Writing the chemical formula of sodium Chloride Sodium is a group 1 element and therefore it has a valency of 1 whilst chlorine is a group 7 element, which also has a valency of 1. Note that there is no need to write 1 as a subscript. ___________________________________________________________________________ Working: Give the Chemical formulae of the following: Calcium Bromide Magnesium Oxide Germanium Oxide ___________________________________________________________________________ Elements with multiple valances: Some elements, particularly the transition metals in the centre block of the periodic table can have more than one valency. Roman numerals are used to show the valency for these elements. We then follow the same set of rules as before. ROMAN NUMERAL VALENCY I 1 II 2 III 3 IV 4 V 5 VI 6 Example: Writing the chemical formula of Copper(II) chloride Copper has a (II) roman numeral, which means that it has a valency of 2. Chlorine is a group 7 element, which has a valency of 1. ___________________________________________________________________________ Working: Give the Chemical formulae of the following: Iron(III)fluoride Manganese(IV)oxide ___________________________________________________________________________ Polyatomic ions: Some compounds contain polyatomic ions i.e. ions that contain more than one kind of atom. Some examples of polyatomic ion are listed below: Charge Name Formula One Positive ammonium NH!" Hydrogencarbonate HCO$ # hydroxide OH $ One Negative nitrate NO$ # permanganate MnO$ ! carbonate CO%$ # Two Negative sulfate SO%$ ! sulfite SO%$ # Three Negative phosphate PO#$ ! The valency of a polyatomic ion is the same as the value of the charge of the ion e.g. NH!" has 1 positive charge so valency = 1 CO%$ # has 2 negative charges so valency = 2 When writing chemical formulae with polyatomic ion always put these ions inside a bracket, e.g. NH!" becomes (NH4) CO%$ # becomes (CO3) Example: Writing the chemical formula of Potassium phosphate Potassium is a group 1 element, which means that it has a valency of 1. Phosphate has three negative charges and thus a valency of 3. Note that the brackets are not needed if there is only one polyatomic ion in the formula so K3(PO4) can be written as K3PO4. ___________________________________________________________________________ Working: Give the Chemical formulae of the following: Magnesium sulfate Calcium nitrate Ammonium hydroxide ___________________________________________________________________________ Names with Prefixes: In the names of some compounds, the ratio of atoms present can be indicated by prefixes, example: Prefix Meaning Mono 1 Di 2 Tri 3 Tetra 4 Penta 5 Hexa 6 As soon as you see a prefix anywhere in the name of a compound do not use the cross valency method and never cancel down ratios. If part of the name of a compound has a prefix whilst another part of the name of the same compound does not have a prefix, assume that only one atom of the element that does not have the prefix is present. Examples include: Carbon monoxide CO Nitrogen dioxide NO2 Phosphorus trichloride PCl3 Phosphorus pentachloride PCl5 Chemical Reactions Substances may interact together and change from one form to another. This happens at the molecular level, where atoms (or groups of atoms) rearrange. This results in bond breaking and formation. These changes are called chemical reactions. Energy is the driving force of all chemical changes, both physical and chemical reactions. Thus, energy is always involved in these reactions. The substances undergoing changes are called reactants, whereas the newly formed substances are called products. Physical appearances of the products can be different from those of the reactants. How do you know that substances are reacting? When a chemical reaction occurs, one of the following usually occurs: Colour change - A colour change indicates that the molecules have changed. Heat content changes - In all chemical reactions, the heat content of the reactants and the heat content of the products is never the same. Sometimes the difference is great and can be easily detected. At other times, the difference is slight and more difficult to detect. Gas produced - Whenever a gaseous product forms in a liquid solution, bubbles can be seen. A colourless gas produced in a reaction of solids is much harder to detect.  Precipitate forms - Precipitates are insoluble products formed by a reaction taking place in a liquid solution. This insoluble product will eventually settle to the bottom, but might immediately appear by turning the clear solution cloudy. Writing Chemical Reactions The most important aspect of a chemical reaction is to know what are the reactants and what are the products. For this, the best description of a reaction is to write an equation for the reaction. A chemical reaction equation gives the reactants and products, and a balanced chemical reaction equation shows the number relationships of reactants and products. Dealing with the quantitative aspect of chemical reactions is called reaction stoichiometry In a chemical reaction we usually also include the ‘state symbols’ to show the physical state of the reactants and products. The state symbols are: (s) – solid (l) – liquid (g) – gas (aq) – aqueous (dissolved in water) Types of Chemical Reactions The vast number of chemical reactions can be classified in any number of ways. One common classification scheme recognizes four major reaction types: combination or synthesis reactions decomposition reactions substitution or single replacement reactions metathesis or double displacement reactions Combination or synthesis reactions: Two or more reactants unite to form a single product. General Reaction Pattern: A + B à AB Examples: S + O2 à SO2 sulphur oxygen sulphur dioxide 2S + 3 O2 à 2 SO3 sulphur oxygen sulphur trioxide 2 Fe + O2 à 2 FeO iron oxygen iron (II) oxide Decomposition reactions: A single reactant is decomposed or broken down into two or more components General Reaction Pattern: AB à A + B Examples: CaCO3 à CaO + CO2 calcium carbonate calcium oxide carbon dioxide 2 H2 O à 2 H2 + O2 water hydrogen oxygen 2 KClO3 à 2 KCl + 3 O2 potassium chlorate potassium chloride oxygen Substitution or Single Replacement Reactions reactions: A single free element replaces or is substituted for one of the elements in a compound. The free element is more reactive than the one its replaces. General Reaction Pattern: A + BC à B + AC Examples: Zn + 2 HCl à H2 + ZnCl2 zinc hydrochloric acid hydrogen zinc chloride Cu + 2 AgNO3 à 2 Ag + Cu(NO3)2 copper silver nitrate silver copper (II) nitrate 2 Na + 2 H2O à 2 NaOH + H2 sodium water sodium hydroxide hydrogen Metathesis or Double Displacement Reactions: This reaction type can be viewed as an "exchange of partners." For ionic compounds, the positive ion in the first compound combines with the negative ion in the second compound, and the positive ion in the second compound combines with the negative ion in the first compound. General Reaction Pattern: AB + CD à AD + CB Examples: HCl + NaOH à NaCl + HOH Hydrochloric acid sodium hydroxide sodium chloride water BaCl2 + 2 AgNO3 à 2 AgCl + Ba(NO3)2 barium chloride silver nitrate silver chloride barium nitrate CaCO3 + 2 HCl à CaCl2 + H2CO3 calcium carbonate hydrochloric acid calcium chloride carbonic acid Balancing of Chemical reactions Even though chemical compounds are broken up and new compounds are formed during a chemical reaction, atoms in the reactants do not disappear nor do new atoms appear to form the products. In chemical reactions, atoms are never created or destroyed. The same atoms that were present in the reactants are present in the products – they are merely reorganized into different arrangements. In a complete chemical equation, the two sides of the equation must be balanced. That is, in a complete chemical equation, the same number of each atom must be present on the reactants and the products sides of the equation. There are two types of numbers that appear in chemical equations. There are subscripts, which are part of the chemical formulas of the reactants and products and there are coefficients that are placed in front of the formulas to indicate how many molecules of that substance is used or produced. The subscripts are part of the formulas and once the formulas for the reactants and products are determined, the subscripts may not be changed. The coefficients indicate the number of each substance involved in the reaction and may be changed in order to balance the equation. Rules for Balancing Chemical Equations 1. Write the correct chemical formulas for all the reactants and the products. 2. Write the formulas of the reactants on the LEFT of the reaction arrow; write the formulas of the products on the RIGHT of the reaction arrow. 3. COUNT the total number of atoms/ions of each element in the reactants and the total number of atoms/ions of each element in the products. A polyatomic ion that appears unchanged on both sides of the equation is counted as a single unit.  4. Balance the elements one at a time using coefficients. When no coefficient is written, the coefficient is assumed to be 1. It is best to begin with elements OTHER THAN hydrogen and oxygen. These elements often occur more than twice in equations. You must NOT attempt to balance the equation by changing subscripts in chemical formulas. 5. Check each atom/ion, or polyatomic ion to be sure that the equation is correctly balanced. 6. Finally, make sure that all the coefficients are in the LOWEST possible whole number ratios. (At least one of the coefficients must be a prime number!) Example: When hydrogen and oxygen react, the product is water. Write the balanced chemical equation. (skeleton) H2 (g) + O2 (g) → H2O (l) 2H 2O | 2H 1O ________________________________________ H2 (g) + O2 (g) → 2 H2O (l) 2H 2O | 4H 2O _________________________________________ BALANCED: 2 H2 (g) + O2 (g) → 2 H2O (l) 4H 2O | 4H 2O ___________________________________________________________________________ Balance the following chemical equations: 1. N2O5 → N2 + O2 2. C5H12 + O2 → CO2 + H2O 3. C8H18 + O2 → CO2 + H2O 4. HC2H3O2 + KNO2 → KC2H3O2 + HNO2 5. AlCl3 + H2O → Al(OH)3 + HCl ___________________________________________________________________________ Catalysts There are various ways that can be used to speed up chemical reactions, for example increasing the reaction temperature increases the rate of reaction. Another way that can be used to speed up chemical reactions is by using Catalysts. These work by providing a different reaction path which has a lower activation energy (the minimum amount of energy that is required for a reaction to start). Although catalysts change the way chemicals react, they are not used up in the chemical reaction i.e. the amount of catalyst at the start of a reaction would be exactly equal to the amount of catalyst at the end of a reaction. For this reason, we do not write the catalyst in the chemical equation. Also, usually small amounts of catalysts are needed to have an effect on a reaction. Equilibrium reactions In most cases, a chemical reaction finishes when all the chemicals you started with (reactants) form the new chemicals (products). However, this is not always the case. In some cases, the products go through a reverse chemical reaction and become the reactants again. So, in this case we will have two chemical reactions which are taking place simultaneously, i.e. Reactants forming the products and the products reforming the reactants. There is a point where you can't tell that any reactions are happening. That's the point when the reaction looks like it is finished. In reality, some of the molecules are turning into products and some are turning back into reactants at the same speed (rate). We say that the reaction is in equilibrium, and thus, these types of reactions are called equilibrium reactions. Le Chatelier’s principle If a system is in equilibrium, and a disturbance is applied to it, the equilibrium will shift in such a direction so as to tend to annul that disturbance. By disturbance here we understand a variation in concentration, pressure or temperature. Acid and bases Acidic and basic are two extremes that describe chemicals, just like hot and cold are two extremes that describe temperature. Mixing acids and bases can cancel out their extreme effects; much like mixing hot and cold water can even out the water temperature. A substance that is neither acidic nor basic is neutral. Definitions of acids and bases There are various theories which are used to describe what acids and bases are. Here we are going to review two of these theories. The Arrhenius Theory This theory was proposed by Arrhenius in 1887. In this theory an acid and a base are defined as follows: An Acid is a substance which was capable of dissociating (the molecules breaks down into ions) in an aqueous solution to produce hydrogen ions (H+). A base is a substance which dissociated in water solution to produce hydroxide ions (OH-). There are several problems with this theory, including: It applies to water-based solutions only Some acids do not contain hydrogens Some bases do not contain hydroxide ions The Bronsted-Lowry Theory This theory was proposed by Bronsted and Lowry in 1923. This approach is generally accepted in biological and medical fields. This theory does not require an aqueous solution or dissociation into ions as in the Arrhenius definition, although it still permits it. In the Bronsted-Lowry theory an acid and a base are defined as follows: An acid is a substance which can donate a hydrogen ion to another substance. A base is a chemical species capable of accepting a proton. When an acid donates a hydrogen ion it forms it conjugate base. A conjugate base has one less hydrogen ion in its molecular formula. Likewise, when a base accepts a hydrogen ion it forms its conjugate acid. A conjugate acid has one extra hydrogen ion in its molecular formula. This idea of conjugate acid-base pairs is an important part of the Bronsted-Lowry approach. When an acid is dissolved in water, it donates a hydrogen ion to the water forming the hydronium ion H3O+. On the other hand, when a base is added to water, the water donates a hydrogen ion to it. Thus, according to this theory, water is acting both a Bronsted-Lowry acid and Bronsted-Lowry base. Thus water is said to be amphoteric (can act both as an acid or a base). When present in water we can classify acid and bases as strong or weak. A strong acid and base will completely dissociate in water or aqueous solution (all the acid or base molecules will donate or accept a hydrogen ion from water respectively). On the other hand, weak acids and bases will only partially dissociate (only some of the acid or base molecules will donate or accept a hydrogen ion from water respectively), forming an equilibrium between acid- conjugate base or the base conjugate acid pairs. The PH scale We can determine how acidic or basic a solution is by the amount of hydronium ions present. Acidic solution will have high hydronium ion concentration while basic solutions will have low hydronium ion concentrations. The accepted way to represent the amount of hydronium ions present is through the mathematical relation called the pH equation. pH = −log[H# O" ] Since the amount of H3O+ is equal to the hydrogen ion, this equation is sometimes written as: pH = −log[H " ] The pH scale ranges from 0 to 14. A pH of 7 is neutral. A pH less than 7 is acidic, and a pH greater than 7 is basic. Each whole pH value below 7 is ten times more acidic than the next higher value. For example, a pH of 4 is ten times more acidic than a pH of 5 and 100 times (10 times 10) more acidic than a pH of 6. The same holds true for pH values above 7, each of which is ten times more basic than the next lower whole value. For example, a pH of 10 is ten times more alkaline than a pH of 9. Pure water is neutral, with a pH of 7.0. When chemicals are mixed with water, the mixture can become either acidic or basic. Highly acidic and basic substances are highly dangerous and can produce severe burns. Common examples of acids and bases Examples of strong acids: hydrochloric acid (HCl) sulfuric acid (H2SO4) Nitric acid (HNO3) Examples of weak acids: ethanoic acid (CH3COOH) carbonic acid (H2CO3) Examples of strong bases: Sodium hydroxide (NaOH) Potassium hydroxide (KOH) Examples of weak bases: Ammonia (NH3) Reactions of Acids and bases When acids react with bases, a salt and water are produced. This reaction is called neutralisation. In general: acid + metal oxide (type of base) → salt + water acid acid + metal hydroxide (type of base) → salt + water Example: NaOH + HCl à NaCl + H2O H2SO4 + Ca(OH)2 à CaSO4 + 2H2O Most bases do not dissolve in water. But if a base can dissolve in water, it is also called an alkali. Note that: A salt is an ionic compound which is made up of two groups of oppositely charged ions. How many of each type of ion the salt has is important because the compound must have an overall electrical charge of zero - that is, an equal balance between positive charge and negative charge. Buffer solutions A buffer solution is one whose pH barely alters on: Adition of small amounts of acid or base On slight dilution A buffer consists of a weak acid and one of its salts or a weak base and one of its salts.