Elements, compounds and mixtures Revision Notes

What is an Empirical Formula?

The empirical formula shows the simplest whole-number ratio of atoms of each element present in a compound. It is not necessarily the same as the molecular formula, which shows the actual number of atoms of each element in a molecule. For example, glucose has a molecular formula of C₆H₁₂O₆, but its empirical formula is CH₂O because the ratio 6:12:6 simplifies to 1:2:1. Understanding this basic ratio is essential for describing and comparing different chemical compounds.

Deducing Empirical Formula from Atom Counts

To find the empirical formula from a given count of atoms, you need to determine the ratio of each type of atom and then simplify it to the smallest whole-number ratio.

1. Count the number of atoms for each element in the compound.
2. Write these counts as a ratio. For example, if a compound has 2 nitrogen atoms and 4 oxygen atoms, the ratio N:O is 2:4.
3. Divide all numbers in the ratio by the greatest common divisor (the largest number that divides into all of them without a remainder). For N:O=2:4, the greatest common divisor is 2, so dividing by 2 gives 1:2.
4. Write the elements with these simplified subscripts. So, the empirical formula of the compound with 2 nitrogen and 4 oxygen atoms is NO₂.

Using Diagrams and Models

When presented with a diagram or model of a compound, the process is similar. Each sphere or representation in the model typically represents an atom of a specific element (often colour-coded). Count the number of each type of atom visible in the diagram or model. Once you have these counts, follow the steps for deducing the empirical formula from atom counts: write the ratio, simplify it by dividing by the greatest common divisor, and then write the empirical formula using the simplified subscripts. This method allows you to visually determine the fundamental composition of a substance.

The Principle of Chromatography

Chromatography separates components of a mixture based on their differential distribution between a stationary phase and a mobile phase. The stationary phase does not move, while the mobile phase carries the sample components through or over the stationary phase. In paper chromatography, paper is the stationary phase and a solvent is the mobile phase. Components that are more soluble in the mobile phase and less adsorbed by the stationary phase travel further.

Identifying Purity Using Chromatography

A pure substance consists of only one chemical compound. When analyzed by chromatography, a pure substance will produce only a single spot on the chromatogram after separation. If a sample is impure, it contains multiple compounds, and each distinct compound will separate and produce its own spot on the chromatogram, resulting in two or more spots. This makes chromatography a quick and effective method for assessing purity.

Performing Paper Chromatography for Purity

To test the purity of a substance using paper chromatography: draw a pencil line near the bottom of a chromatography paper. Place a small spot of the substance onto this line. Suspend the paper in a beaker with a suitable solvent (the mobile phase) so that the solvent level is below the pencil line. Allow the solvent to run up the paper. Once the solvent front is near the top, remove the paper and mark the solvent front. If visible spots appear, observe how many there are. If only one spot is present, the substance is pure; if multiple spots appear, it is impure.

Interpreting Chromatograms for Purity

After drying the chromatogram (often needing a locating agent for colourless substances), examine the number of separated spots. A single spot confirms purity. Multiple spots indicate an impure sample, as each spot represents a different chemical component. The distance a component travels relative to the solvent front (Rf value) is characteristic of that component under specific conditions and helps in identification, but for purity, the key is simply the number of spots observed.

Atomic Structure and Electron Shells

Every atom consists of a nucleus containing protons and neutrons, surrounded by electrons in electron shells. The atomic number represents the number of protons in an atom's nucleus and, for a neutral atom, also the number of electrons. Electrons fill shells starting from the innermost shell. The first shell can hold up to 2 electrons, the second up to 8, and the third (for the purpose of stability in common reactions) up to 8. The outermost shell contains valence electrons, which are key to determining an atom's chemical reactivity.

Reactivity and Electron Configuration

Atoms react to achieve a stable electron configuration, usually a full outer electron shell. This often involves achieving the electron configuration of a noble gas. Elements with an incomplete outer shell will participate in chemical reactions to gain, lose, or share electrons. The number of valence electrons dictates how an atom will achieve this stability. For example, atoms with one, two, or three valence electrons tend to lose them to form positive ions, while atoms with five, six, or seven valence electrons tend to gain or share electrons to form negative ions or covalent bonds, respectively.

Groups and Reactivity Trends

The position of an element in the Periodic Table, specifically its group number, directly relates to the number of valence electrons. Group 1 elements (alkali metals) have one valence electron and are highly reactive as they readily lose this electron. Group 17 elements (halogens) have seven valence electrons and are also highly reactive, needing to gain just one electron. Group 18 elements (noble gases) have a full outer shell (8 valence electrons, or 2 for helium) and are therefore chemically unreactive, as they already possess a stable electron configuration.

Atomic Number and Chemical Identity

The atomic number uniquely identifies an element because it determines the number of protons. In a neutral atom, the atomic number also sets the total number of electrons, which then dictates how these electrons are arranged in shells. This electron arrangement, especially the number of electrons in the outermost shell, is the fundamental reason behind an element's specific chemical properties and reactivity. Thus, different atomic numbers lead to different electron arrangements and different chemical behaviours.

Physical Properties of Metals

Metals generally possess a set of distinctive physical properties. They are typically lustrous and shiny due to their ability to reflect light. Most metals are solid at room temperature, with the notable exception of mercury, which is a liquid. They are excellent conductors of both heat and electricity, a property attributed to the delocalised electrons within their metallic structure. Furthermore, metals are malleable, meaning they can be hammered or pressed into different shapes without breaking, and ductile, which allows them to be drawn into wires. They also tend to have high melting and boiling points, and relatively high densities.

Chemical Properties of Metals

From a chemical perspective, metals are characterised by their tendency to lose electrons from their outer shell during chemical reactions. This loss of electrons results in the formation of positive ions (cations). This electron-losing tendency makes metals reactive, with more reactive metals losing electrons more readily. When metals react with oxygen, they typically form metal oxides, which are often basic in nature. For example, sodium reacts with oxygen to form sodium oxide (Na₂O). Some metals also react with acids to produce hydrogen gas and a salt.

Physical Properties of Non-Metals

Non-metals display a broader range of physical properties compared to metals. At room temperature, non-metals can exist as solids (e.g., carbon, sulfur), liquids (e.g., bromine), or gases (e.g., oxygen, nitrogen). They are generally dull in appearance and, if solid, are often brittle and break easily rather than bending or stretching. Unlike metals, non-metals are poor conductors of heat and electricity; they are good insulators. An important exception is graphite, a form of carbon, which is an excellent electrical conductor. Non-metals typically have lower melting and boiling points and lower densities than metals.

Chemical Properties of Non-Metals

Chemically, non-metals tend to gain or share electrons to achieve a stable outer electron shell. When they gain electrons, they form negative ions (anions). When they share electrons, they form covalent bonds, leading to molecules or giant covalent structures. Non-metals react with oxygen to form non-metal oxides, which are typically acidic (e.g., carbon dioxide) or neutral (e.g., carbon monoxide). Many non-metals also react with metals, forming ionic compounds where the non-metal acts as the electron acceptor.