Classification of Matter, Atomic Structure, and Energy Systems

Comprehensive practice flashcards covering classification of matter, properties and uses of elements, compounds, alloys, atomic structure, historical models of the atom, periodic table organization, physical versus chemical change, chemical equations, cellular metabolism, and energy transfer modes.

What is an element, and what are some common examples found in everyday objects?

An element is a pure substance made up of only one type of atom that cannot be broken down into simpler substances by chemical means. Common examples include:

• Iron ($\text{Fe}$): used in steel support structures, tools, and cookware.
• Aluminium ($\text{Al}$): used in drink cans, food wrap foil, and aircraft.
• Copper ($\text{Cu}$): used in electrical wiring and plumbing pipes.
• Carbon ($\text{C}$): found as graphite in pencils and charcoal.
• Oxygen ($\text{O}$): atmospheric oxygen gas ($\text{O}_2$) in the air.
• Silicon ($\text{Si}$): used in glass, ceramics, and microchips.

What are the orthographic rules for writing chemical symbols of elements?

Each element has a unique standardized one- or two-letter symbol. The first letter is always strictly capitalised, and the second letter, when present, is strictly lowercase (e.g., $\text{C}$, $\text{Al}$, $\text{Fe}$).

How do the physical and chemical properties of metals differ from non-metals?

Metals:

• Located on the left side and centre of the periodic table.
• Lustrous (shiny), malleable (can be hammered into sheets), and ductile (can be drawn into wires).
• Solid at room temperature ($20\,^{\circ}\text{C}$ to $25\,^{\circ}\text{C}$), with mercury ($\text{Hg}$) as the sole liquid exception.
• Excellent conductors of heat and electricity; high density and melting points.
• Low electronegativity; readily lose valence electrons to form positive cations and react with acids to produce hydrogen gas ($\text{H}_2$).

Non-metals:

• Located on the upper right side of the periodic table (plus hydrogen on the top left).
• Dull and non-reflective; brittle solids that crumble when bent or stressed.
• Exist in all three states at room temperature: gases (e.g., $\text{O}_2, \text{N}_2, \text{Cl}_2$), a liquid ($\text{Br}_2$), and solids (e.g., $\text{C}, \text{S}, \text{P}$).
• Poor conductors (insulators, except graphite); low melting and boiling points.
• High electronegativity; tend to gain or share electrons to form anions or covalent compounds; show no reaction with dilute acids.

Which elements are the sole liquid exceptions at standard room temperature ($20\,^{\circ}\text{C}$ to $25\,^{\circ}\text{C}$)?

Mercury ($\text{Hg}$) is the only metal that is liquid at room temperature, and bromine ($\text{Br}_2$) is the only non-metal that is liquid at room temperature.

What are metalloids, and what six elements belong to this category?

Metalloids (semimetals) sit along the diagonal 'staircase' dividing line separating metals from non-metals. They possess a mixture of metallic and non-metallic properties: typically solid at room temperature, shiny in appearance, but mechanically brittle. They function as semiconductors (conducting electricity moderately well under specific conditions, such as elevated temperature or applied voltage) and exhibit chemical behaviors predominantly like non-metals.

The complete list of metalloids is:

1. Boron ($\text{B}$)
2. Silicon ($\text{Si}$)
3. Germanium ($\text{Ge}$)
4. Arsenic ($\text{As}$)
5. Antimony ($\text{Sb}$)
6. Tellurium ($\text{Te}$)

How do the specific properties of copper, aluminium, and iron relate to their primary uses?

• Copper ($\text{Cu}$): High electrical conductivity and ductility make it ideal for electrical cables, wiring, and plumbing pipes.
• Aluminium ($\text{Al}$): Low density (lightweight), high malleability, and strong corrosion resistance (due to a protective oxide layer) make it useful for commercial aircraft fuselages, drink cans, and foil wraps.
• Iron ($\text{Fe}$): Massive structural rigidity and high tensile strength when alloyed with carbon into steel make it ideal for structural beams in bridges, skyscrapers, heavy machinery, and cookware.

How do the specific properties of gold, silicon, and stainless steel determine their practical applications?

• Gold ($\text{Au}$): Exceptional chemical inertness (does not corrode or tarnish), extreme malleability, and conductivity make it valuable for fine jewellery and corrosion-free electronic connectors.
• Silicon ($\text{Si}$): Semiconductor capability allows precise control of electrical current, making it essential for microprocessors, computer chips, solar cells, and transistors.
• Stainless Steel: Strong resistance to rusting and corrosion combined with exceptional hardness and tensile strength makes it ideal for cutlery, sinks, surgical instruments, and food-processing equipment.

What is the key scientific distinction between a compound and an alloy?

• Compound: A pure substance composed of two or more distinct elements chemically bonded together in an invariant, fixed stoichiometric ratio (e.g., water, $\text{H}_2\text{O}$, in a fixed $2:1$ ratio; sodium chloride, $\text{NaCl}$, in a fixed $1:1$ ionic lattice).
• Alloy: A solid physical mixture containing a parent metal fused with one or more other elements (e.g., bronze, brass, steel). An alloy is a mixture, NOT a compound; its components are not chemically bonded in a fixed stoichiometric ratio.

What are the elemental compositions and properties of stainless steel, bronze, brass, and carbon steel?

• Stainless Steel: Iron ($\text{Fe}$), chromium ($\text{Cr}$), nickel ($\text{Ni}$), and carbon ($\text{C}$). Highly resistant to rust and corrosion via chromium oxide passivation.
• Bronze: Copper ($\text{Cu}$) and tin ($\text{Sn}$). Significantly harder and more wear-resistant than pure copper, ductile, and resistant to saltwater corrosion (used in bearings, marine hardware, sculptures).
• Brass: Copper ($\text{Cu}$) and zinc ($\text{Zn}$). High acoustic resonance, low friction, malleable, and golden appearance (used in musical instruments, door hardware).
• Carbon Steel: Iron ($\text{Fe}$) with interstitial carbon ($\text{C}$). Greatly increases tensile strength and hardness over pure iron (used in bridge beams, rebar, railway tracks).

Why is the atom defined as the smallest unit of an element that retains its properties?

An atom retains all unique chemical properties and identity of that element (e.g., a single copper atom is still copper). If an atom is disassembled further into its subatomic constituent particles (protons, neutrons, electrons), those isolated fragments completely lose the identity and chemical behavior of the original element.

Identify the labelled subatomic particles and architectural regions of the atom shown in this diagram.

• Nucleus: The dense central core containing protons and neutrons.
• Proton: Positively charged nucleon located in the nucleus.
• Neutron: Neutral nucleon located in the nucleus.
• Electron: Negatively charged particle orbiting the nucleus.
• Electron Shell: Concentric energy level surrounding the nucleus occupied by electrons.

What are the location, relative electrical charge, and relative mass of protons, neutrons, and electrons?

• Proton: Location = Nucleus; Relative Charge = $+1$; Relative Mass = $1\,\text{amu}$.
• Neutron: Location = Nucleus; Relative Charge = $0$; Relative Mass = $1\,\text{amu}$.
• Electron: Location = Electron shells / extranuclear cloud; Relative Charge = $-1$; Relative Mass = $\frac{1}{1836}\,\text{amu}$.

Note: Almost the entire mass of an atom ($> 99.9\%$) is concentrated in the dense nucleus.

What are Atomic Number ($Z$) and Mass Number ($A$), and how do they relate to subatomic particle counts in a neutral atom?

• Atomic Number ($Z$): The total number of protons in the nucleus of an atom. It defines the elemental identity.
• In any neutral atom: $\text{Number of Protons} = \text{Number of Electrons}$.
• Mass Number ($A$): The total combined count of nucleons ($\text{Protons} + \text{Neutrons}$) in the nucleus.
• Neutron Formula: $\text{Number of Neutrons} = \text{Mass Number} - \text{Atomic Number}$.

An atom of sodium ($\text{Na}$) has an Atomic Number of $11$ and a Mass Number of $23$. How many protons, electrons, and neutrons does it possess?

• Protons = $11$ (equal to the atomic number)
• Electrons = $11$ (equal to the protons in a neutral atom)
• Neutrons = $23 - 11 = 12$ ($\text{Mass Number} - \text{Atomic Number}$)

What were the atomic hypotheses proposed by Democritus and John Dalton?

• Democritus ($500\,\text{B.C.}$): Proposed a philosophical deduction that all matter is composed of minute, indivisible, indestructible fundamental particles called 'atomos' (meaning uncuttable).
• John Dalton ($1800\text{s}$): Formulated the first modern scientific atomic theory (Solid Sphere / Billiard Ball Model), proposing that atoms are solid, indivisible spheres that combine in whole-number ratios to form compounds and rearrange during reactions.

What experimental breakthrough did J.J. Thomson achieve in $1897$, and what atomic model did he propose?

Using cathode ray tube deflection experiments, J.J. Thomson discovered the first subatomic particle: the negatively charged electron. He proposed the 'Plum Pudding Model', picturing the atom as a diffuse spherical cloud of positive charge embedded with discrete negative electrons, like raisins in plum pudding.

How did Ernest Rutherford's $1909$ gold foil experiment debunk the Plum Pudding Model?

Rutherford fired alpha particles at ultra-thin gold foil. Most passed straight through, but a small fraction deflected at extreme angles or bounced backward. This disproved the diffuse Plum Pudding Model and led to the Nuclear Model, showing that the atom is mostly empty space with virtually all positive charge and mass concentrated in an extraordinarily tiny central nucleus.

What modification did Niels Bohr introduce to the atomic model in $1913$?

To resolve the instability of Rutherford's orbiting electrons, Bohr proposed the Planetary / Shell Model. Electrons move exclusively in fixed, quantized, circular orbits (electron shells) at defined radii from the nucleus without emitting radiation, and only jump between shells by absorbing or emitting specific quanta of energy.

When and by whom were the proton and neutron discovered?

• Proton: Discovered in $1919$ by Ernest Rutherford through artificial transmutation experiments bombarding nitrogen gas with alpha particles.
• Neutron: Discovered in $1932$ by James Chadwick by bombarding beryllium atoms with alpha particles, detecting neutral radiation composed of particles with roughly the mass of a proton.

What are the administrative and structural details of the Year 8 Assessment Task No. 2 at Smith's Hill High School?

• Institution: Smith's Hill High School
• Task & Focus: Year 8 Assessment Notification Task No. 2; Periodic table and atomic structure; Change
• Date & Time: Tuesday, 3rd November, Period 2, running from $9:45\,\text{am}$ to $11:05\,\text{am}$
• Location: Multi-Purpose Unit (MPU)
• Total Marks: $35\,\text{marks}$
• Time Allocated: $1\,\text{hour}$ working time plus $5\,\text{minutes}$ reading time
• Structure: Section I ($20\,\text{multiple-choice questions} = 20\,\text{marks}$); Section II (short-answer and extended response questions = $15\,\text{marks}$)
• Syllabus Outcomes: SC4-PRT-01, SC4-CHG-01, SC4-WS-03, SC4-WS-05

What are the query protocols for the Smith's Hill High School Year 8 Assessment Task?

• Verbal in-class questions are prohibited to preserve equity.
• Questions must be submitted digitally via email at least $2\,\text{days}$ prior to the assessment date; answers are posted universally via Google Classroom.
• Any post-assessment marking query must be written directly on the test front page for teacher review.

How is the modern periodic table organized in terms of periods and groups?

• Periods: $7$ horizontal rows. Moving left to right across a period, atomic number increases by $1$ proton per element, and the number of occupied electron shells remains constant. Properties transition systematically from reactive metals to metalloids, non-metals, and finally unreactive noble gases.
• Groups: Vertical columns. Elements in the same group possess the same number of outer-shell valence electrons and consequently exhibit analogous chemical behaviors, valencies, and reaction patterns.

What are the key chemical and physical traits of Group 1, Group 2, Group 17, and Group 18 elements?

• Group 1 (Alkali Metals): $1$ valence electron; extremely reactive (react vigorously with water to produce alkaline hydroxides and $\text{H}_2$ gas); soft, low density.
• Group 2 (Alkaline Earth Metals): $2$ valence electrons; highly reactive metals forming $+2$ cations and basic oxides.
• Group 17 (Halogens): $7$ valence electrons; highly reactive non-metals forming corrosive diatomic molecules; readily gain $1$ electron to form $-1$ halide anions.
• Group 18 (Noble Gases): Completely filled valence shell ($8$ electrons octet, or $2$ duet for Helium); chemically inert, monoatomic, colorless, odorless gases.

Based on this periodic table tile for Nobelium, identify what each indicated element of the tile represents.

• Top Number ($102$): Atomic number (number of protons in the nucleus).
• Central Symbol ($\text{No}$): Chemical symbol.
• Bottom Text (nobelium): Element name.
• Upper-right Number in Parentheses ($(259)$): Atomic weight (mass number of the most stable isotope).
• Lower Formula ($[\text{Rn}]5f^{14}7s^2$): Electron configuration.
• Lower-right Symbol / Bars: Physical state indicator (solid at $20\,^{\circ}\text{C}$ / $68\,^{\circ}\text{F}$).

Identify the element name, symbol, and atomic number for: (1) Metal in Group 2, Period 4; (2) Metalloid in Group 14, Period 3; (3) Metalloid in Group 15, Period 4.

1. Metal in Group 2, Period 4: Calcium ($\text{Ca}$, Atomic number $20$)
2. Metalloid in Group 14, Period 3: Silicon ($\text{Si}$, Atomic number $14$)
3. Metalloid in Group 15, Period 4: Arsenic ($\text{As}$, Atomic number $33$)

What are the maximum electron capacities of the first three electron shells, and what are the electron configurations of Hydrogen, Helium, Neon, and Argon?

Maximum Capacities:

• Shell 1 ($n = 1$): $2$ electrons
• Shell 2 ($n = 2$): $8$ electrons
• Shell 3 ($n = 3$): $8$ electrons (for the first 18 elements)

Configurations:

• Hydrogen ($\text{H}, Z = 1$): $1$
• Helium ($\text{He}, Z = 2$): $2$ (first shell fully closed)
• Neon ($\text{Ne}, Z = 10$): $2,8$ (second shell closed octet)
• Argon ($\text{Ar}, Z = 18$): $2,8,8$ (third shell closed octet)

What are the key differences and observational indicators distinguishing physical change from chemical change?

• Physical Change: Involves changes in state of matter (melting, evaporation), morphology (size, shape, texture), or dissolution without altering chemical composition. Readily reversible by physical means.
• Chemical Change (Chemical Reaction): Involves breaking and reforming chemical bonds to synthesize entirely new substances with different properties. Typically difficult or impossible to reverse by physical means.

Observational Indicators of Chemical Change:

1. Persistent, unexpected change in colour
2. Noticeable change or production of scent/odour
3. Spontaneous production of gas (fizzing, bubbling, effervescence)
4. Formation of an insoluble solid precipitate from two solutions
5. Enthalpy/temperature change (heating or cooling) or emission of light

What observations and word equations describe the reaction of heating iron and sulfur, and the reaction between vinegar and bicarbonate of soda?

1. Thermal Synthesis of Iron Sulfide:

• Before: Grey magnetic iron powder ($\text{Fe}$) and bright yellow non-magnetic sulfur powder ($\text{S}$) exist as a physically separable mixture.
• After heating: Vigorous exothermic glow forms iron sulfide ($\text{FeS}$), a uniform black, non-magnetic, brittle solid that cannot be separated with a magnet.
• Word Equation: $\text{Iron} + \text{Sulfur} \rightarrow \text{Iron Sulfide}$

1. Vinegar and Bicarbonate of Soda:

• Before: Clear pungent liquid (acetic acid) and fine white crystalline powder (sodium hydrogen carbonate).
• After: Violent effervescence releasing carbon dioxide gas; temperature drops (endothermic process); white powder dissolves.
• Word Equation: $\text{Vinegar} + \text{Bi-carb Soda} \rightarrow \text{Carbon Dioxide} + \text{Water} + \text{Sodium Acetate}$

What is the difference between exothermic and endothermic reactions, and how do photosynthesis and cellular respiration exemplify them?

• Exothermic Reaction: Releases net thermal energy to the surroundings, causing a measurable temperature increase (e.g., combustion, iron-sulfur synthesis, cellular respiration).
• Endothermic Reaction: Absorbs net thermal energy from the surroundings, causing a measurable temperature decrease (e.g., vinegar and baking soda, photosynthesis).

Biochemical Examples:

• Photosynthesis (Endothermic): Absorbs solar radiant energy inside chloroplasts to synthesize glucose.   $\text{Carbon Dioxide} + \text{Water} \rightarrow \text{Glucose} + \text{Oxygen}$
• Cellular Respiration (Exothermic): Oxidizes glucose inside mitochondria to liberate stored energy as ATP and heat.   $\text{Glucose} + \text{Oxygen} \rightarrow \text{Carbon Dioxide} + \text{Water}$

What are the three modes of thermal energy transfer?

1. Conduction: Thermal transfer through direct physical contact where vibrating particles collide and pass kinetic energy along a thermal gradient (most effective in dense solids and metals with free electrons).
2. Convection: Thermal transfer through bulk displacement of heated fluids (liquids and gases). Warmer fluid expands, becomes less dense, and rises, while cooler, denser fluid sinks, forming a convection current.
3. Radiation: Direct transfer of energy via electromagnetic waves (chiefly infrared). Requires no medium and travels through a complete vacuum at the speed of light ($3.00 \times 10^8\,\text{m/s}$).

What are the formulas and definitions for Kinetic Energy, Gravitational Potential Energy, and Chemical Potential Energy?

• Kinetic Energy ($\text{KE}$): Energy possessed by an object due to its motion.
  $\text{KE} = \frac{1}{2}mv^2$
• Gravitational Potential Energy ($\text{GPE}$): Energy stored in an object elevated against a gravitational field.
  $\text{GPE} = mgh$
• Chemical Potential Energy: Energy stored within the stable chemical bonds holding atoms and molecules together, released or absorbed when bonds break and reform during reactions.

What is the thermodynamic difference between an open system and a closed system?

• Open System: Freely exchanges both energy and physical matter across its boundary with the external environment (e.g., an unsealed boiling beaker of water losing heat and steam mass).
• Closed System: Physically sealed boundary prevents any exchange of matter, permitting solely the exchange of energy (heat, work, radiation) into and out of the system (e.g., a sealed rigid metal canister submerged in water).

What does the Law of Conservation of Energy state, and what is an example of an energy transformation?

The Law of Conservation of Energy states that energy can neither be created nor destroyed in any physical or chemical process; it can only be transferred from one place to another or transformed from one form into another. The total energy in an isolated system remains invariant.

Examples:

• A falling object transforms gravitational potential energy ($\text{GPE}$) directly into kinetic energy ($\text{KE}$).
• Photosynthesis transforms radiant electromagnetic energy from the Sun into stored chemical bond energy within glucose molecules.
• An engine transforms chemical potential energy of petrol into mechanical kinetic energy along with waste thermal and acoustic energy.