OCR GCSE Combined Science A Gateway revision notes
Predicting and identifying reactions and products Revision Notes
Study Predicting and identifying reactions and products with curriculum-aligned Revision Notes resources, practice links, and exam-focused support.
At a glance
revision notes
Resource type
Topic
Predicting and identifying reactions and products
Revision notes
Establishing an Order of Reactivity for Metals Revision Notes
Understanding Metal Reactivity
Metal reactivity refers to the ease with which a metal loses electrons to form positive ions, leading to chemical reactions. Highly reactive metals readily react with substances like oxygen, water, and acids, whereas less reactive metals react more slowly or not at all. The reactivity of metals determines their position in the reactivity series.
Investigating Reactivity with Water
One way to establish an order of reactivity is by observing how metals react with water. Very reactive metals (e.g., potassium, sodium, calcium) react vigorously with cold water, producing hydrogen gas and a metal hydroxide. Moderately reactive metals (e.g., magnesium, zinc, iron) may react slowly with cold water or only with steam. Less reactive metals (e.g., lead, copper, silver) show no reaction with water or steam. The vigour of the reaction indicates relative reactivity.
Investigating Reactivity via Displacement Reactions
A key method to determine metal reactivity is through displacement reactions. A more reactive metal will displace a less reactive metal from a solution of its salt. For example, if metal A displaces metal B from a solution of metal B's salt, then metal A is more reactive than metal B. If no reaction occurs, metal A is less reactive than metal B. This allows for a direct comparison of reactivity between two metals.
Experimental Procedure: Metal Reactivity
To compare metals, you could place small pieces of each metal into separate test tubes containing dilute acid (e.g., hydrochloric acid), cold water, or steam. For displacement, place a metal strip into a salt solution of another metal. Observe the vigour of gas production (bubbles), temperature changes, or the formation of a new solid (for displacement). A faster, more vigorous reaction indicates higher reactivity. Use controlled variables like metal surface area and solution concentration to ensure fair comparisons.
Predicting Reactivity from the Periodic Table Revision Notes
Understanding Groups and Periods
The Periodic Table arranges elements by atomic number. Elements in the same vertical column (group) have similar chemical properties because they possess the same number of outer shell electrons, which dictate how they react. Elements in the same horizontal row (period) have the same number of electron shells, but their chemical properties change gradually across the period.
Predicting Metal Reactivity
Metals typically react by losing electrons to form positive ions. For metals, reactivity generally increases as you move down a group. This is because the outer electrons are further from the nucleus and are screened by more inner electron shells, making them easier to remove. Reactivity decreases as you move from left to right across a period because the increased nuclear charge pulls the outer electrons more strongly, making them harder to lose. Group 1 metals are the most reactive metals.
Predicting Non-metal Reactivity
Non-metals typically react by gaining or sharing electrons. For non-metals, reactivity generally decreases as you move down a group. This is because the atomic radius increases, and the outer electrons are further from the positive nucleus, reducing the attraction for incoming electrons. Reactivity increases as you move from left to right across a period because the increased nuclear charge and smaller atomic radius enhance the attraction for incoming electrons. Group 17 elements (halogens) are the most reactive non-metals.
Applying Reactivity Trends to Reactions
Knowing the reactivity trends allows us to predict the outcomes of reactions. For example, a more reactive metal will displace a less reactive metal from its compound. Similarly, a more reactive halogen will displace a less reactive halogen from its metal halide solution. These predictions are based directly on the element's position within the Periodic Table.
Metal Reactivity and Ion Formation Revision Notes
Understanding Metal Reactivity
Metal reactivity describes how readily a metal undergoes chemical reactions. At its core, this involves the metal atoms losing electrons to become positively charged ions (cations). The easier a metal loses electrons, the more reactive it is considered. For example, sodium reacts vigorously with water because its atoms readily shed an electron. This fundamental tendency to form positive ions dictates its position in the reactivity series.
Reactions with Water and Dilute Acids
Metals react with water or dilute acids by losing electrons to hydrogen ions, which then form hydrogen gas. In reactions with water, active metals can displace hydrogen from water molecules, producing a metal hydroxide and hydrogen gas. With dilute acids, metals react to displace hydrogen from the acid, forming a metal salt and hydrogen gas. The vigour of these reactions is a direct measure of the metal's reactivity – more reactive metals react more enthusiastically. For instance, potassium reacts explosively with water, while magnesium reacts steadily with dilute hydrochloric acid.
The Link to Ion Formation
The 'tendency to form its positive ion' refers to how easily a metal atom loses its outer shell electrons. Metals are inherently electropositive, meaning they have a propensity to lose electrons. A highly reactive metal has a strong tendency to lose electrons and thus readily forms its positive ion. Conversely, a less reactive metal holds onto its electrons more tightly, making it harder to form its positive ion. This directly explains why reactive metals sit higher in the reactivity series and engage in displacement reactions, as they are strong enough to 'take' electrons from less reactive metals' ions.
Practical Investigations (PAG C4, C5)
PAGs (Practical Activity Groups) C4 and C5 often involve experiments to observe and compare the reactivity of different metals. This typically includes reacting metals with water (cold and steam) and dilute acids (e.g., hydrochloric acid). By observing the rate of hydrogen gas production, the temperature change, and the extent of the reaction, students can deduce the relative reactivity of metals. For instance, metals that produce hydrogen bubbles rapidly and evolve significant heat are more reactive, demonstrating their greater tendency to form positive ions compared to metals that react slowly or not at all.
Properties of Groups 1, 7, and 0 elements and their relation to outer shell electrons Revision Notes
Group 1: Alkali Metals (1 outer electron)
Group 1 elements all have one electron in their outer shell, making them highly reactive. They lose this electron easily to form positive ions (+1 charge). Reactivity increases down the group because the outer electron is further from the nucleus and more shielded, so it is lost more easily. Physical properties like melting and boiling points decrease down the group due to weaker metallic bonding as atoms get larger and the forces of attraction between the positive ions and delocalised electrons weaken. They are soft metals, good conductors of electricity and heat, and have a low density.
Group 7: Halogens (7 outer electrons)
Group 7 elements have seven electrons in their outer shell and tend to gain one electron to form negative ions (-1 charge) or share electrons. They exist as diatomic molecules (e.g., Cl₂). Reactivity decreases down the group because the outer shell is further from the nucleus and more shielded, making it harder to attract an electron to complete the outer shell. Melting and boiling points increase down the group because the size of the molecules increases, leading to stronger intermolecular forces requiring more energy to overcome.
Group 0: Noble Gases (Full outer shell)
Group 0 elements have a full outer shell of electrons (8 electrons, except for helium with 2). This stable electron configuration means they are very unreactive (inert) and do not readily form compounds. They exist as single atoms (monatomic). They have very low melting and boiling points due to weak intermolecular forces between the atoms. Melting and boiling points increase slightly down the group because the larger atoms have more electrons, leading to slightly stronger temporary intermolecular forces (London dispersion forces).
Predicting Trends based on Outer Electrons
The number of outer electrons determines an element's group and primary chemical behaviour. For Group 1 metals, the ease of losing the outer electron dictates reactivity. For Group 7 non-metals, the ease of gaining an electron dictates reactivity. For Group 0 noble gases, the full outer shell dictates their inertness. General trends within a group (e.g., increasing atomic size and shielding down a group) explain variations in reactivity, melting points, and boiling points.
Related topics
Study nearby topics next
Revision hubs