Year 9 and 10 Science Practice Flashcards

Comprehensive practice flashcards covering Year 9 and 10 Science topics: Atomic structure, periodic table trends, ionic bonding, response to stimuli, nervous and endocrine systems, reproduction, energy, carbon cycle, waves, sound, light, disease, and space exploration.

Atom

The basic building block of all matter, consisting of a central nucleus containing protons and neutrons, surrounded by electrons in electron shells.

Proton

A positively charged subatomic particle with a charge of $+1$ and a relative mass of $1\,\text{unit}$, located in the nucleus of all atoms.

Neutron

A neutral subatomic particle with no electrical charge ($0$) and a mass of $1\,\text{unit}$, located in the nucleus of an atom.

Electron

A negatively charged subatomic particle with a charge of $-1$ and a mass of approximately $\frac{1}{2000}$ of a proton, occupying most of the volume outside the nucleus.

Nucleus

The dense, positively charged central region of an atom that contains protons and neutrons (nucleons).

Nucleon

A collective term for the subatomic particles located within an atomic nucleus, specifically protons and neutrons.

Atomic Number ($Z$)

The number of protons found in the nucleus of an atom, which uniquely determines the element's identity and its chemical properties.

Mass Number ($A$)

The total number of protons and neutrons (nucleons) present within the nucleus of a specific atom.

Nuclide Symbol

A standard chemical notation representing an individual atom in the format ${}_{Z}^{A}\text{X}$, where $A$ is the mass number, $Z$ is the atomic number, and $\text{X}$ is the element symbol.

Isotopes

Atoms of the same element that possess identical atomic numbers (same number of protons) but different mass numbers due to differing numbers of neutrons.

Carbon-12

A stable isotope of carbon with $6$ protons and $6$ neutrons, having a mass number of $12$.

Carbon-13

A stable isotope of carbon with $6$ protons and $7$ neutrons, having a mass number of $13$.

Carbon-14

A radioisotope of carbon with $6$ protons and $8$ neutrons, possessing a mass number of $14$ and a half-life of $5600\,\text{years}$.

Nuclear Decay

The spontaneous process by which unstable atomic nuclei break apart to form new, more stable nuclei, releasing nuclear radiation.

Radioisotope

An unstable isotope of an element that undergoes nuclear decay and emits nuclear radiation.

Alpha Particle ($\alpha$)

A fast-moving helium nucleus containing two protons and two neutrons (${}_{2}^{4}\text{He}$) that has a charge of $+2$.

Beta Particle ($\beta$)

A highly energised electron or positron emitted from an unstable nucleus with no mass and either a negative or positive charge.

Gamma Ray ($\gamma$)

A very high-energy photon of electromagnetic radiation with no mass and no electrical charge, emitted during nuclear decay.

Penetrating Power of Alpha Radiation

Low penetration ability; alpha particles cannot travel easily through materials and can be stopped by a sheet of paper or human skin.

Penetrating Power of Beta Radiation

Moderate penetration ability; beta particles can penetrate human skin and damage living tissue, but can be stopped by thin layers of plastic, wood, or aluminium.

Penetrating Power of Gamma Radiation

High penetration ability; gamma rays cause serious damage to living tissue and can only be effectively stopped by thick shields of lead or concrete.

Half-Life

The time required for half of the radioactive nuclei in a given sample of a radioisotope to undergo nuclear decay.

Potassium-40

A radioisotope of potassium used to date rocks, having a half-life of approximately $1.25\,\text{billion years}$.

Radium-226

A radioactive isotope with a half-life of $1600\,\text{years}$, taking $3200\,\text{years}$ for $75\%$ of its sample nuclei to decay.

Plutonium-236

A radioisotope whose sample decays exponentially over time, exhibiting a half-life of approximately $2.8\,\text{years}$.

Electron Shell

A designated region of space surrounding the atomic nucleus where electrons are arranged in order of increasing energy levels.

First Electron Shell Capacity

The innermost electron shell of an atom, which can accommodate a maximum of $2$ electrons.

Second Electron Shell Capacity

The second electron shell from the nucleus, which can accommodate a maximum of $8$ electrons.

Third Electron Shell Capacity (Elements $1$–$20$)

For the first $20$ elements, the third electron shell can hold a maximum of $8$ electrons before electrons begin filling the fourth shell.

Fourth Electron Shell Capacity (Elements $1$–$20$)

For the first $20$ elements, the fourth electron shell accommodates a maximum of $2$ electrons (as seen in potassium and calcium).

Electron Configuration

The numerical representation of how electrons are distributed among the electron shells of an atom, written from innermost shell to outermost.

Electron Configuration of Chlorine-35

The shell electron arrangement of a neutral chlorine atom with $17$ electrons, written as $2,8,7$.

Valence Shell

The outermost occupied electron shell of an atom that determines how the atom chemically interacts and bonds with other atoms.

Valence Electrons

The electrons located in the outermost shell (valence shell) of an atom that participate in chemical bonding.

Octet Rule

The chemical principle that atoms tend to gain, lose, or share electrons until their valence shell is stable with $8$ electrons (or $2$ for the first shell).

Periodic Table

An arrangement of all known chemical elements in order of increasing atomic number, reflecting recurring periodic trends in chemical and physical properties.

Period

A horizontal row in the periodic table; elements in the same period have the same number of occupied electron shells.

Group

A vertical column in the periodic table (numbered $1$ to $18$); elements in the same group possess the same number of valence electrons and share similar chemical properties.

Alkali Metals

The soft, highly reactive metallic elements located in Group $1$ of the periodic table, each possessing $1$ valence electron.

Alkaline Earth Metals

The metallic elements found in Group $2$ of the periodic table, which each possess $2$ valence electrons.

Halogens

The reactive non-metallic elements in Group $17$ of the periodic table that possess $7$ valence electrons and form diatomic molecules.

Noble Gases

The odourless, colourless, unreactive gaseous non-metals in Group $18$ of the periodic table that have full valence shells.

Periodic Trend: Atomic Size Down a Group

The atomic radius increases down a group because each successive element adds an extra occupied electron shell.

Periodic Trend: Atomic Size Across a Period

The atomic radius decreases from left to right across a period as increasing nuclear charge pulls the electron shells closer.

Periodic Trend: Metal Reactivity Down a Group

The reactivity of metals increases down a group because the single valence electron is further from the nucleus and more easily lost.

Periodic Trend: Non-Metal Reactivity Down a Group

The reactivity of non-metals decreases down a group because the incoming electron experiences less attractive force from the nucleus.

Ion

An atom or molecule that has gained or lost one or more valence electrons, acquiring an overall net positive or negative electrical charge.

Cation

A positively charged ion formed when an atom (typically a metal) loses electrons, resulting in more protons than electrons.

Anion

A negatively charged ion formed when an atom (typically a non-metal) gains electrons, resulting in more electrons than protons.

Electrostatic Attraction

The force of electrical attraction acting between oppositely charged particles, such as positive cations and negative anions.

Ionic Bonding

A chemical bond formed by the electrostatic attraction between oppositely charged ions arranged in a regular crystalline lattice.

Ionic Lattice

A three-dimensional repeating crystalline network of alternating cations and anions held rigidly together by electrostatic attraction.

Overall Charge of an Ionic Compound

The total electrical charge of an ionic compound, which is always neutral ($0$) because positive and negative charges balance completely.

Monatomic Anion Naming Rule

The naming convention where the ending of a non-metal element's name is dropped and replaced with the suffix '-ide' (e.g. chlorine becomes chloride).

Fixed-Charge Cations

Metal cations that form only one stable ionic charge, including Group $1$ ($+1$), Group $2$ ($+2$), aluminium ($\text{Al}^{3+}$), zinc ($\text{Zn}^{2+}$), and silver ($\text{Ag}^+$).

Variable-Charge Cations

Transition metals that can form ions with multiple possible charges, indicated in their chemical names using Roman numerals (e.g. iron(II) as $\text{Fe}^{2+}$ and iron(III) as $\text{Fe}^{3+}$).

Polyatomic Ion

A charged chemical species composed of two or more covalently bonded atoms that acts as a single ionic unit.

Ammonium Ion

A positively charged polyatomic cation with the formula $\text{NH}_4^+$ and a charge of $+1$, formed by ammonia reacting with a hydrogen ion.

Hydroxide Ion

A negatively charged polyatomic anion with the chemical formula $\text{OH}^-$ and a charge of $-1$.

Nitrate Ion

An oxygen-containing polyatomic anion with the chemical formula $\text{NO}_3^-$ and an electrical charge of $-1$.

Carbonate Ion

A polyatomic anion with the chemical formula $\text{CO}_3^{2-}$ and an electrical charge of $-2$.

Sulfate Ion

A polyatomic anion containing sulfur and oxygen, having the formula $\text{SO}_4^{2-}$ and an electrical charge of $-2$.

Phosphate Ion

A polyatomic anion with the chemical formula $\text{PO}_4^{3-}$ and an electrical charge of $-3$.

Sulfide vs Sulfate

Sulfide is the single-atom anion $\text{S}^{2-}$, whereas sulfate is the polyatomic oxygen-containing anion $\text{SO}_4^{2-}$.

Nitride vs Nitrate

Nitride is the monatomic anion $\text{N}^{3-}$, while nitrate is the polyatomic anion $\text{NO}_3^-$.

Magnesium Nitrate Formula

The neutral ionic formula $\text{Mg(NO}_3\text{)}_2$, requiring brackets to indicate two nitrate anions balance one $\text{Mg}^{2+}$ cation.

Calcium Carbonate Formula

The neutral ionic formula $\text{CaCO}_3$, formed by combining one $\text{Ca}^{2+}$ ion with one $\text{CO}_3^{2-}$ ion in a $1:1$ ratio.

Aluminium Oxide Formula

The neutral ionic formula $\text{Al}_2\text{O}_3$, formed by balancing two $\text{Al}^{3+}$ cations with three $\text{O}^{2-}$ anions.

Brittleness of Ionic Compounds

The mechanical property where an applied force shifts lattice layers, aligning like charges that repel and cause the crystal to shatter.

Melting Points of Ionic Compounds

Typically very high temperatures because strong electrostatic forces between alternating cations and anions require substantial thermal energy to overcome.

Solid-State Electrical Conductivity of Ionic Compounds

Ionic solids do not conduct electricity because the charged ions are fixed rigidly in lattice positions and cannot move freely.

Aqueous Electrical Conductivity of Ionic Compounds

Ionic compounds dissolved in water conduct electricity because their dissociated cations and anions are free to move and carry charge.

Aqueous Solution

A homogeneous liquid mixture in which water serves as the solvent that dissolves solutes.

Solute

A substance that is dissolved within another substance (the solvent) to form a solution.

Solvent

The liquid substance in which a solute is dissolved to form a homogeneous solution.

Precipitate

An insoluble solid product that separates from an aqueous solution when two compatible soluble ionic solutions are mixed.

SNAPE Rule

A chemical mnemonic stating that ionic compounds containing Sodium ($\text{Na}^+$), Nitrate ($\text{NO}_3^-$), Ammonium ($\text{NH}_4^+$), Potassium ($\text{K}^+$), or Ethanoate ($\text{CH}_3\text{COO}^-$) are always soluble.

CHOPS Rule

A mnemonic identifying ions that are typically insoluble in water—Carbonate ($\text{CO}_3^{2-}$), Hydroxide ($\text{OH}^-$), Oxide ($\text{O}^{2-}$), Phosphate ($\text{PO}_4^{3-}$), and Sulfide ($\text{S}^{2-}$)—unless paired with a SNAPE cation.

Flame Test

An analytical qualitative procedure used to identify specific metal cations based on the characteristic colour of light they emit when heated in a Bunsen burner flame.

Sodium Flame Test Colour

A persistent bright yellow-orange flame produced when sodium cations are heated.

Strontium Flame Test Colour

A distinctive bright red flame produced when strontium cations are introduced into a flame.

Copper(II) Flame Test Colour

A characteristic blue-green flame observed when copper(II) ions are heated.

Hydroxide Test

A chemical test adding sodium hydroxide to an aqueous solution to identify cations based on the specific colour and appearance of the insoluble precipitate formed.

Iron(III) Hydroxide Test Result

The formation of a characteristic red-brown precipitate when sodium hydroxide solution is mixed with an orange $\text{Fe}^{3+}$ solution.

Iron(II) Hydroxide Test Result

The formation of a green precipitate when sodium hydroxide solution is added to a yellow $\text{Fe}^{2+}$ solution.

Manganese(II) Hydroxide Test Result

The formation of a pale brown precipitate in a nearly colourless solution upon adding sodium hydroxide to an $\text{Mn}^{2+}$ solution.

Magnesium Hydroxide Test Result

The appearance of a white precipitate in a colourless solution when sodium hydroxide is mixed with a solution containing $\text{Mg}^{2+}$.

Ammonium Test with Litmus Paper

A test where adding sodium hydroxide and warming generates alkaline ammonia ($\text{NH}_3$) gas that turns moist red litmus paper blue.

Carbonate Anion Acid Test

A test where adding dilute hydrochloric acid produces bubbles and fizzing due to the rapid release of carbon dioxide ($\text{CO}_2$) gas.

Sulfate Test with Barium Chloride

A confirmatory test where adding barium chloride to an aqueous sulfate solution forms a dense, insoluble white precipitate of barium sulfate ($\text{BaSO}_4$).

Hierarchical Organisation of Multicellular Organisms

The structural biological levels of complexity progressing from cells to tissues, organs, organ systems, and finally the complete organism.

Stimulus

A detectable physical or chemical change in the internal or external environment of an organism that is registered by a sensory receptor.

External Stimulus

A change originating outside the organism's body, such as environmental temperature, light, sound, or physical contact.

Internal Stimulus

A physiological change originating inside the organism's body, such as fluctuations in blood glucose, blood pressure, or internal body temperature.

Receptor

A specialised sensory cell or organ that detects a specific stimulus and converts it into a biological signal.

Response

A physiological action, movement, or behavioural change executed by an effector in reaction to a detected stimulus.

Central Nervous System (CNS)

The primary coordinating centre of the human nervous system, comprising the brain and the spinal cord.

Peripheral Nervous System (PNS)

The network of nerves outside the central nervous system that links sensory receptors and effectors to the brain and spinal cord.

Spinal Cord

A column of nerve fibres extending from the base of the brain down the back that conducts signals to and from the brain and controls spinal reflexes.

Neuron

A specialised nerve cell designed to transmit electrical and chemical messages throughout the nervous system.