Formulas Revision Notes

A chemical formula is a symbolic representation that shows which elements are present in a substance and how many atoms of each element are combined. It uses chemical symbols for elements and numbers called subscripts to indicate the quantity of atoms.

For example, the formula H2O indicates a molecule made of two hydrogen atoms and one oxygen atom. The subscript '2' after hydrogen shows there are two hydrogen atoms, while the absence of a subscript after oxygen means there is one oxygen atom.

• Chemical symbols represent elements (e.g., H for hydrogen, O for oxygen).
• Subscripts indicate the number of atoms of each element in the molecule.
• If no subscript is shown, it means one atom of that element.

Elements can exist as single atoms or molecules made of atoms bonded together. The chemical formula for an element depends on its natural form.

For example, noble gases like helium exist as single atoms, so helium's formula is He. However, some elements like oxygen naturally exist as molecules of two atoms bonded together, so oxygen's formula is O2.

This distinction is important because the formula must reflect the actual form of the element under normal conditions.

• Single atom elements have formulas with one symbol and no subscript (e.g., He).
• Diatomic elements have formulas with a subscript 2 (e.g., O2, N2, H2).

Compounds are substances formed when atoms of different elements chemically combine in fixed ratios. Their chemical formulas show the types of atoms and their ratios.

For example, sodium chloride (table salt) has the formula NaCl, meaning one sodium atom combines with one chlorine atom. Water has the formula H2O, meaning two hydrogen atoms combine with one oxygen atom.

The formula reflects the simplest whole-number ratio of atoms in the compound, which helps identify the compound and predict its properties.

• Compounds contain two or more different elements.
• Formulas show the ratio of atoms in the compound (e.g., NaCl is 1:1, H2O is 2:1).
• Subscripts indicate the number of atoms of each element in the compound.

Knowing chemical formulas allows chemists to calculate the amounts of substances involved in reactions and to write balanced chemical equations.

For example, in water (H2O), the formula tells us that for every molecule, there are two hydrogen atoms and one oxygen atom. This helps in calculating the mass of hydrogen and oxygen needed to produce water.

Understanding formulas also helps in identifying substances and predicting their behaviour in chemical reactions.

• Formulas help calculate relative masses of elements in compounds.
• They assist in balancing chemical equations by showing atom ratios.
• They are essential for identifying substances in chemical analysis.

Here are some common elements and compounds with their chemical formulas:

• Hydrogen gas: H2
• Oxygen gas: O2
• Nitrogen gas: N2
• Carbon dioxide: CO2
• Methane: CH4
• Sodium chloride: NaCl
• Glucose: C6H12O6

A molecular formula is a symbolic representation that shows the exact number and types of atoms present in a single molecule of a compound. Unlike empirical formulas, which show the simplest ratio of atoms, molecular formulas give the actual count of each atom in one molecule.

For example, the molecular formula for water is H2O. This tells us that each water molecule contains two hydrogen atoms and one oxygen atom. The subscript numbers indicate the quantity of each atom type; if there is no subscript, it means there is only one atom of that element in the molecule.

• Molecular formula shows actual numbers of atoms in a molecule.
• Subscripts indicate the number of atoms of each element.
• No subscript means one atom of that element.

The letters in a molecular formula represent the chemical symbols of the elements present in the molecule. Each symbol corresponds to a specific type of atom. For example, in carbon dioxide, CO2, 'C' stands for carbon atoms and 'O' stands for oxygen atoms.

The order of elements in a molecular formula usually follows conventions, such as listing carbon first, hydrogen second, and then other elements alphabetically, but the key point is that each symbol identifies the atom type included in the molecule.

• Chemical symbols represent atom types.
• Each symbol corresponds to one element.
• Order of symbols follows chemical conventions but does not affect meaning.

The subscript numbers following each element symbol indicate how many atoms of that element are present in one molecule. For example, in methane, CH4, the '4' after hydrogen means there are four hydrogen atoms attached to one carbon atom.

If there is no subscript after an element symbol, it means there is only one atom of that element in the molecule. For example, in ammonia, NH3, there is one nitrogen atom and three hydrogen atoms.

• Subscripts show the number of atoms of each element.
• No subscript means one atom.
• Example: H2O has two hydrogen atoms and one oxygen atom.

Molecular formulas are essential for understanding the composition of compounds. They allow chemists to know exactly what atoms and how many of each are present in a molecule, which is crucial for predicting chemical behaviour and reactions.

Knowing the molecular formula helps in calculating molecular mass, balancing chemical equations, and understanding the properties of substances.

• Shows exact composition of molecules.
• Helps calculate molecular mass.
• Essential for balancing chemical equations and predicting reactions.

Here are some simple examples illustrating molecular formulas:

Water: H2O (2 hydrogen atoms, 1 oxygen atom) Carbon dioxide: CO2 (1 carbon atom, 2 oxygen atoms) Methane: CH4 (1 carbon atom, 4 hydrogen atoms) Oxygen gas: O2 (2 oxygen atoms)

These examples show how molecular formulas clearly communicate the composition of molecules.

• Water: H2O Carbon dioxide: CO2 Methane: CH4 Oxygen gas: O2

Molecular models and diagrams visually represent molecules by showing the atoms and their connections. Each atom is represented by a sphere or symbol, often colour-coded or labelled to indicate the element. For example, hydrogen atoms might be white spheres and oxygen atoms red spheres in a ball-and-stick model. Diagrams may show element symbols connected by lines representing bonds.

Recognising the elements present is the first step. Each distinct type of atom corresponds to a chemical element, identified by its symbol (e.g., H for hydrogen, O for oxygen).

• Models use spheres or symbols to represent atoms.
• Colours or labels help identify elements.
• Lines or sticks represent chemical bonds between atoms.

After identifying the elements, count how many atoms of each element appear in the model or diagram. This count must be exact because the chemical formula depends on the number of atoms of each element in one molecule.

For example, if a model shows two hydrogen atoms and one oxygen atom connected, the counts are: hydrogen = 2, oxygen = 1.

• Count each atom type carefully.
• Ignore bonds when counting atoms; focus on the spheres or symbols.
• Ensure the count corresponds to one molecule, not multiple molecules.

The chemical formula is written by placing the element symbols side by side, followed by subscripts indicating the number of atoms of each element. If there is only one atom of an element, no subscript is written.

Using the previous example, with 2 hydrogen atoms and 1 oxygen atom, the formula is H₂O. The subscript 2 shows two hydrogen atoms; oxygen has no subscript because there is only one atom.

The order of elements in the formula usually follows conventions: hydrogen is often written first if present, followed by oxygen, or elements are ordered according to their electronegativity or group in the periodic table for more complex compounds.

• Write element symbols in correct order.
• Use subscripts to show the number of atoms.
• Omit the subscript if there is only one atom of that element.

Example 1: A model shows 1 carbon atom and 4 hydrogen atoms. Counting gives carbon = 1, hydrogen = 4. The formula is CH₄ (methane).

Example 2: A diagram shows 2 nitrogen atoms bonded together. Counting gives nitrogen = 2. The formula is N₂ (nitrogen gas).

Example 3: A model shows 1 carbon atom, 2 oxygen atoms. Counting gives carbon = 1, oxygen = 2. The formula is CO₂ (carbon dioxide).

• CH₄: 1 carbon, 4 hydrogen atoms.
• N₂: 2 nitrogen atoms.
• CO₂: 1 carbon, 2 oxygen atoms.

A frequent error is miscounting atoms, especially when atoms are hidden behind others in a 3D model. Rotate the model or carefully examine the diagram to ensure all atoms are counted.

Another mistake is writing a subscript of 1, which is incorrect; no subscript should be written for a single atom.

Confusing ionic formulas with molecular formulas can lead to errors; this objective focuses only on simple molecular compounds, where atoms are covalently bonded in discrete molecules.

• Check all atoms are counted, including those partially obscured.
• Do not write subscript 1; omit it instead.
• Focus on covalent molecules, not ionic compounds.

A word equation is a way to represent a chemical reaction using the names of the substances involved rather than their chemical formulas. It shows what substances react together (reactants) and what new substances are formed (products). This method is useful for visualising chemical changes, especially when chemical formulas are not yet known or when focusing on the overall process rather than detailed atomic composition.

• Reactants are substances present before the reaction starts.
• Products are substances formed as a result of the reaction.
• An arrow (→) separates reactants and products, indicating the direction of the reaction.

In a word equation, the reactants are listed on the left side of the arrow, and the products are listed on the right side. The arrow points from reactants to products, showing the transformation during the reaction. Multiple reactants or products are separated by plus signs (+).

For example, in the reaction where hydrogen gas reacts with oxygen gas to form water, the word equation is:

Hydrogen + Oxygen → Water. This shows that hydrogen and oxygen are the starting materials, and water is the substance produced.

• Reactants → Products
• Use plus signs (+) to separate multiple reactants or products.
• The arrow (→) means 'react to form' or 'produce'.

To write a word equation, first identify the substances involved before and after the reaction. Then place the reactants on the left, products on the right, and connect them with an arrow. If there is more than one reactant or product, separate them with plus signs.

For example, when calcium carbonate decomposes on heating to form calcium oxide and carbon dioxide, the word equation is:

Calcium carbonate → Calcium oxide + Carbon dioxide.

• Identify reactants and products from the reaction description.
• Write reactants on the left, products on the right.
• Use plus signs for multiple substances on either side.

Word equations help to understand what substances are involved in a reaction and how they change. By reading a word equation, you can tell which substances are consumed and which are formed. This is useful for predicting products of reactions and for communicating chemical changes clearly.

For example, the word equation: Sodium + Chlorine → Sodium chloride shows that sodium and chlorine react to form sodium chloride, a compound with different properties from the reactants.

• Reactants are consumed during the reaction.
• Products are newly formed substances.
• Word equations summarise the chemical change without showing atomic details.

Practice writing word equations for common reactions such as combustion, decomposition, and displacement. For example, the combustion of methane can be written as: Methane + Oxygen → Carbon dioxide + Water.

In exams, ensure the correct names of substances are used and the arrow is placed correctly. Avoid mixing chemical formulas in word equations unless specifically asked. Check that all reactants and products mentioned in the question are included.

Remember, word equations do not show quantities or balance atoms; that is done in symbol equations.

• Use correct chemical names, not formulas, in word equations.
• Include all reactants and products mentioned.
• Place reactants on the left, products on the right, separated by an arrow.
• Do not balance atoms in word equations; that is for symbol equations.

In any chemical reaction, atoms are neither created nor destroyed; they are simply rearranged. This means the total number of atoms of each element must be the same before and after the reaction. A chemical equation represents this reaction using chemical formulas for reactants and products. Balancing the equation ensures it obeys this conservation of atoms.

For example, in the reaction between hydrogen and oxygen to form water, the number of hydrogen and oxygen atoms must be equal on both sides of the equation.

• Atoms are conserved in chemical reactions.
• Balancing equations reflects this conservation.
• Chemical formulas show the types and numbers of atoms in substances.

Balancing a chemical equation involves placing numbers called coefficients in front of formulas to adjust the quantity of molecules or moles. It is important not to change the subscripts within formulas, as these define the chemical identity of the substances.

For example, the unbalanced equation for hydrogen reacting with oxygen is: H2 + O2 → H2O. Counting atoms shows 2 hydrogen atoms on the left but 2 on the right, and 2 oxygen atoms on the left but only 1 on the right. To balance oxygen, place a coefficient 2 before H2O: H2 + O2 → 2H2O. Now hydrogen atoms are 2 on the left and 4 on the right, so place 2 before H2: 2H2 + O2 → 2H2O. This balances both elements.

• Coefficients multiply the entire formula to balance atoms.
• Never change subscripts within formulas.
• Balance one element at a time, usually starting with metals or complex molecules.

State symbols indicate the physical state of each substance in a chemical equation: (s) for solid, (l) for liquid, (g) for gas, and (aq) for aqueous solution (dissolved in water). Including these symbols provides additional information about the reaction conditions and helps interpret the reaction more fully.

For example, the balanced equation for the combustion of methane is: CH4(g) + 2O2(g) → CO2(g) + 2H2O(l). This shows methane and oxygen are gases, carbon dioxide is a gas, and water is liquid under the reaction conditions.

• State symbols clarify physical states of substances.
• They are placed in parentheses after each formula.
• Common symbols: (s), (l), (g), (aq).

1. Write the correct formulas for all reactants and products, including state symbols if known.

2. Count the number of atoms of each element on both sides.

3. Choose an element to balance first, usually one that appears in only one reactant and one product to simplify the process, often metals or complex molecules first, then hydrogen and oxygen last as they often appear in multiple compounds or as diatomic molecules (O2, H2). 4. Adjust coefficients to balance that element without changing subscripts. 5. Repeat for all elements until all are balanced. 6. Check the final equation to confirm the same number of atoms of each element on both sides and that coefficients are in the simplest whole number ratio.

• Write formulas and state symbols first.
• Balance elements one at a time using coefficients.
• Check atom counts after each adjustment.
• Simplify coefficients if possible.

Once balanced, chemical equations can be analyzed to determine the relative amounts of reactants and products involved in a reaction. The coefficients represent mole ratios, which can be used to calculate masses, volumes (for gases), or numbers of particles.

For example, in the balanced equation 2H2 + O2 → 2H2O, the coefficient 2 before H2 means 2 moles of hydrogen react with 1 mole of oxygen to produce 2 moles of water. This ratio is essential for stoichiometric calculations in quantitative chemistry.

• Coefficients indicate mole ratios.
• Use balanced equations for quantitative calculations.
• State symbols help predict reaction conditions and phases.

Ionic compounds consist of positively charged ions (cations) and negatively charged ions (anions) held together by electrostatic forces. Because the compound is electrically neutral, the total positive charge must exactly balance the total negative charge. This principle is fundamental when determining the formula of an ionic compound.

For example, if a compound contains sodium ions (Na⁺) and chloride ions (Cl⁻), the charges are +1 and -1 respectively. Since one Na⁺ balances one Cl⁻, the formula is NaCl. If the charges are not equal, the number of ions must be adjusted to balance the total charge.

• Ionic compounds are neutral overall.
• Total positive charge = total negative charge.
• Formulas reflect the ratio of ions needed for charge neutrality.

To find the formula, first identify the charges on the ions involved. For many common ions, charges are predictable: alkali metals form +1 ions, alkaline earth metals +2, halogens -1, and oxygen usually -2. Transition metals may have variable charges and require given information.

Knowing the charge on each ion allows you to calculate how many of each ion are needed to balance the charges.

• Use periodic table groups to predict ion charges.
• Common cations: Na⁺, Mg²⁺, Al³⁺.
• Common anions: Cl⁻, O²⁻, N³⁻.
• Variable charges require specific information.

Once the charges are known, determine the smallest whole-number ratio of ions that results in zero net charge. Multiply the ion charges by the number of ions to ensure total positive and negative charges are equal.

For example, aluminium ions (Al³⁺) and oxide ions (O²⁻):

Al³⁺ has a charge of +3 per ion, O²⁻ has -2 per ion. To balance, find the least common multiple of 3 and 2, which is 6. This means 2 Al³⁺ ions (2 × +3 = +6) and 3 O²⁻ ions (3 × -2 = -6) balance charges, so the formula is Al₂O₃.

• Calculate total positive and negative charges.
• Adjust ion numbers to balance charges.
• Use the simplest whole-number ratio.

Models or diagrams often show ions as spheres with different colours or labels. The relative number of each ion shown corresponds to the ratio in the formula. Count the ions and write the formula using these ratios.

For example, a model showing 1 calcium ion (Ca²⁺) and 2 chloride ions (Cl⁻) indicates the formula CaCl₂, since the charges balance (+2 and 2 × -1 = -2).

Always check that the charges balance when using models.

• Count ions shown in the model.
• Write formula using these ratios.
• Confirm charge neutrality.

Example 1: Determine the formula of a compound formed from magnesium ions (Mg²⁺) and chloride ions (Cl⁻).

Mg²⁺ charge: +2; Cl⁻ charge: -1. To balance charges, 1 Mg²⁺ ion pairs with 2 Cl⁻ ions.

Formula: MgCl₂. Example 2: A model shows 1 aluminium ion (Al³⁺) and 3 oxide ions (O²⁻). Verify the formula and charge balance: 1 × +3 = +3; 3 × -2 = -6, charges do not balance. Adjust to 2 Al³⁺ and 3 O²⁻ ions: 2 × +3 = +6; 3 × -2 = -6, balanced. Formula: Al₂O₃.

• Apply charge balance method.
• Check simplest ratio.
• Confirm neutrality.

Chemical symbol equations represent chemical reactions using chemical formulas and symbols to show the substances involved. Each formula shows the elements and their proportions in a compound or element.

State symbols indicate the physical state of each substance in the reaction: (s) for solid, (l) for liquid, (g) for gas, and (aq) for aqueous solution (dissolved in water). Including state symbols clarifies the conditions of the reaction and helps predict reaction behaviour.

For example, the reaction of solid zinc with hydrochloric acid solution is written as: Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g). Here, zinc is solid, hydrochloric acid and zinc chloride are aqueous, and hydrogen is a gas.

• State symbols: (s) solid, (l) liquid, (g) gas, (aq) aqueous solution.
• They provide information about physical states and reaction conditions.
• Always include state symbols in symbol equations for clarity.

A balanced chemical equation has the same number of each type of atom on both sides, obeying the law of conservation of mass. Balancing ensures the equation accurately represents the reaction.

To balance an equation, adjust coefficients (whole numbers before formulas) to equalise atom counts on reactants and products sides. Never change subscripts within formulas as that changes the substances.

Example: The reaction of hydrogen gas with oxygen gas to form water is unbalanced as H2 + O2 → H2O. Balancing gives 2H2 + O2 → 2H2O, meaning 2 molecules of hydrogen react with 1 molecule of oxygen to produce 2 molecules of water.

• Count atoms of each element on both sides.
• Adjust coefficients to balance atoms.
• Check final equation for lowest whole number coefficients.

Ionic equations show the ions involved in reactions occurring in aqueous solutions, focusing on species that change during the reaction.

In aqueous solutions, many compounds dissociate into ions. For example, sodium chloride dissolves as Na+(aq) and Cl−(aq). Writing ionic equations involves splitting soluble ionic compounds into their ions.

Example: The reaction of sodium hydroxide solution with hydrochloric acid solution forms sodium chloride and water. The full ionic equation is: Na+(aq) + OH−(aq) + H+(aq) + Cl−(aq) → Na+(aq) + Cl−(aq) + H2O(l).

• Split soluble ionic compounds into ions in aqueous solution.
• Show all ions present before and after reaction.
• Include state symbols to indicate physical states.

Spectator ions are ions present on both sides of the ionic equation that do not participate in the chemical change. Removing them simplifies the equation to the net ionic equation, showing only the species involved in the reaction.

In the previous example, Na+(aq) and Cl−(aq) are spectator ions. Removing them gives the net ionic equation: OH−(aq) + H+(aq) → H2O(l). This equation highlights the actual chemical change: acid-base neutralisation.

Writing net ionic equations helps focus on the essential chemistry and is often required in exams.

• Identify ions unchanged on both sides (spectator ions).
• Remove spectator ions to write the net ionic equation.
• Net ionic equations show only reacting species.

Always write chemical formulas correctly, using correct symbols and subscripts.

Include state symbols to gain full marks and show understanding of reaction conditions.

Balance equations carefully; check atom counts after balancing to avoid errors in coefficients or subscripts mistakes (which change substances). Use the smallest whole number coefficients possible to simplify equations for exams and clarity in communication. When writing ionic equations, ensure only aqueous ionic compounds are split into ions; solids, liquids, and gases remain as molecules or formula units in the equation. Practice writing full ionic and net ionic equations from word equations or symbol equations to build confidence.

• Use correct chemical formulas and symbols.
• Include state symbols for all substances.
• Balance equations with smallest whole number coefficients.
• Split only aqueous ionic compounds into ions for ionic equations.
• Remove spectator ions to write net ionic equations.