UNIT 1 aos1 Flashcards
(29 cards)
what is the order of electron configuration?
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d
7s 7p
s=2
p=6
d=10
f=14
what is the atomic radius and its trend
measurement for the size of atoms; distance from nucleus to outermost shells
increases down a group (more electron shells)
decreases across a period (shells pulled tightly together due to effective nuclear charge)
what is electronegativity and its trend
Ability of an element to attract electrons towards itself when forming a chemical bond
Increase across periods
Decrease down groups
Atomic Size: Smaller atoms have higher electronegativity because the nucleus is closer to the bonding electrons.
Nuclear Charge: More protons in the nucleus increase the ability to attract electrons.
Shielding Effect: More inner electrons reduce the pull of the nucleus on outer electrons, lowering electronegativity.
Electron Configuration: Atoms closer to completing their outer shells have a stronger desire for electrons.
what is ionization energy and its trend
Ionisation energy is the amount of energy required to remove one electron from a neutral atom in its gaseous state
Increases across a period
Decreases down a group
Atomic Size: Larger atoms have lower ionisation energy because their outer electrons are farther from the nucleus and experience a weaker pull.
Nuclear Charge: A greater number of protons in the nucleus increases ionisation energy, as the stronger positive charge holds the electrons more tightly.
Shielding Effect: When inner electrons block the pull of the nucleus on outer electrons, the ionisation energy decreases. This effect becomes more pronounced in atoms with many inner electron shells.
Electron Configuration: Atoms with stable electron configurations (like noble gases) have much higher ionisation energies because removing an electron disrupts their stability.
what is metal reactivity and its trend
Metal reactivity refers to how easily a metal can undergo chemical reactions, particularly with water, acids, or oxygen.
Increases down a group
Decreases across a period
Ionisation Energy: Lower ionisation energy means metals lose electrons more easily, increasing reactivity.
Atomic Size: Larger atoms have weaker nuclear pull on outer electrons, making them more reactive.
Shielding Effect: More inner electron shells reduce the effective nuclear charge, making it easier for the metal to lose electrons.
Electron Configuration: Metals close to achieving a stable electron configuration by losing electrons (e.g., alkali metals) are highly reactive.
what is nuclear effective charge and its trend
Effective nuclear charge is the net positive charge experienced by an electron in an atom. It reflects the balance between the attractive pull of the protons in the nucleus and the repelling effect of inner electrons
Formula:
𝑍eff=𝑍−𝑆, where: 𝑍 is the total number of protons (atomic number).
𝑆 is the shielding constant, accounting for inner electrons blocking nuclear pull.
Decreases down a group
Increases across a period
what are covalent bonds?
Covalent bonds are a type of chemical bond that occurs when two or more atoms share electrons.
what is the VSEPR theory?
based on the principle that electron groups around an atom repel each other- consequently they are arranged as far as possible from each
- can help predict shapes of simple molecules
define element
An element is a pure substance made up of only one type of atom. It cannot be broken down into simpler substances by chemical means.
define isotope
An isotope is a variant of a chemical element that has the same number of protons but a different number of neutrons in its nucleus. This means isotopes of an element have the same atomic number (which determines the element) but different mass numbers due to the variation in neutrons.
define ion
An ion is an atom or molecule that has gained or lost one or more electrons, resulting in a net electric charge.
what are critical elements?
On the periodic table, critical elements often refer to chemical elements essential for specific industries, technologies, or biological processes.
Conserves Resources: Extends the availability of finite elements like lithium and rare earths.
Protects Environment: Reduces mining, habitat destruction, and greenhouse gas emissions.
Reduces Waste: Prevents valuable materials from ending up in landfills.
Economic Benefits: Lowers costs and creates jobs in sustainable industries.
polar vs non polar molecules
Polar Molecules: Have an uneven charge distribution due to differences in electronegativity and asymmetrical molecular shapes, leading to a net dipole moment (e.g., H₂O).
Non-Polar Molecules: Have an even charge distribution because either the bonds are non-polar or the molecule is symmetrical, so dipole moments cancel out (e.g., CO₂).
intramolecular bond vs intermolecular force
Intramolecular bonds are the strong chemical bonds that hold atoms together within a molecule. eg. covalent bond, ionic bond, metallic bond
An intermolecular force is the attraction or repulsion that occurs between molecules, affecting their physical properties such as boiling and melting points.
types of intermolecular forces
Dispersion Forces (London Forces):
Present in all molecules, regardless of whether they are polar or nonpolar.
Result from temporary fluctuations in electron distribution, creating momentary dipoles.
Generally weak but can become stronger in larger, heavier molecules.
Dipole-Dipole Forces:
Occur between molecules that have permanent dipoles (polar molecules).
The partially positive end of one molecule is attracted to the partially negative end of another.
Hydrogen Bonding:
A stronger type of dipole-dipole force.
Arises when a hydrogen atom bonded to a highly electronegative atom (like oxygen, nitrogen, or fluorine) interacts with a lone pair of electrons on another electronegative atom.
properties of covalent substances
Low Melting and Boiling Points:
Covalent substances typically have lower melting and boiling points compared to ionic compounds. This is because the intermolecular forces (such as Van der Waals forces or hydrogen bonds) holding the molecules together are weaker than the strong electrostatic forces in ionic compounds.
Poor Electrical Conductivity:
Covalent compounds do not conduct electricity in solid or liquid form because they do not have free-moving charged particles (such as ions or electrons) like ionic compounds or metals.
Solubility:
Many covalent compounds are soluble in non-polar solvents (like benzene or hexane), but they are often insoluble in polar solvents like water. However, some covalent compounds (such as sugar) can dissolve in polar solvents because they form hydrogen bonds with the solvent molecules.
structure and bonding of diamond
Structure and Bonding:
Atomic Arrangement: Diamond has a tetrahedral structure, where each carbon atom is covalently bonded to four other carbon atoms in a 3D network.
Covalent Bonding: The carbon atoms form strong, directional covalent bonds with each other, creating an incredibly strong and rigid three-dimensional lattice.
Bonding Overview: The strong covalent bonds make the diamond structure very stable and durable.
2. Properties and Explanation:
Hardness:
Diamond is the hardest known material because of the strong covalent bonds extending in all directions throughout the structure. This makes it difficult to break the bonds and deform the material. This hardness makes diamond useful in cutting, grinding, and drilling tools.
Electrical Conductivity:
Diamond does not conduct electricity because there are no free electrons or ions to move. All of its electrons are tightly held in covalent bonds, and there is no “delocalization” of charge carriers. As a result, diamond is an insulator.
Heat Conductivity:
Diamond has excellent thermal conductivity. The strong covalent bonds in diamond allow the efficient transfer of vibrational energy (heat) through the lattice. In fact, diamond is one of the best conductors of heat due to the strong bonding and the tight, regular arrangement of atoms that facilitates rapid heat transfer.
Transparency:
Diamond is also transparent and is used in high-end optical applications (e.g., lenses) because it allows light to pass through without significant absorption.
3. Applications of Diamond:
Cutting Tools: Diamond is used in tools for cutting, grinding, and drilling, especially in industries that require extreme durability.
Electronics: Despite being an insulator, diamond can be used in electronics for heat dissipation due to its high thermal conductivity.
Jewelry: Its hardness and brilliance make diamond a prized material in the jewelry industry.
structure and bonding of graphite
Structure and Bonding:
Atomic Arrangement: Graphite has a layered structure, with each layer consisting of carbon atoms bonded to three other carbon atoms in a hexagonal arrangement. The layers are held together by weak van der Waals forces, allowing them to slide past one another easily.
Bonding Overview: Within each layer, the carbon atoms are bonded through strong covalent bonds. Between the layers, the forces are weak, allowing the layers to slide over each other.
2. Properties and Explanation:
Hardness:
Graphite is soft and slippery because the weak van der Waals forces between the layers allow the layers to slide over each other. This makes graphite useful as a lubricant and in pencil leads.
Electrical Conductivity:
Graphite conducts electricity. This is because one electron from each carbon atom in the hexagonal lattice is free to move between layers, creating a “sea” of delocalized electrons. These electrons can move within the layers, making graphite an excellent conductor of electricity along the plane of the layers. However, graphite does not conduct electricity vertically between layers.
Heat Conductivity:
Graphite is also a good conductor of heat along the planes of the layers. The strong covalent bonds within the layers allow heat to be transferred efficiently. However, like its electrical conductivity, heat conductivity is much lower perpendicular to the layers due to the weak forces between them.
Malleability:
Due to the weak forces between layers, graphite is flexible and can be molded or shaped in different forms.
3. Applications of Graphite:
Lubricants: Graphite is used as a lubricant in applications where high temperatures are present and conventional oils break down.
Pencils: Graphite’s soft, slippery nature makes it perfect for writing and drawing (as in pencil leads).
Electrical Conductivity: Graphite is used in electrodes, batteries, and electrical brushes because of its ability to conduct electricity.
Heat Shields: Its ability to withstand high temperatures and conduct heat makes graphite useful in high-temperature applications, such as heat shields and refractory materials.
define allotrope
An allotrope is a different form of the same element in the same physical state, where the atoms are bonded together in different ways, leading to different physical and chemical properties.
properties of metals
Lustre (Shininess)
Explanation: Metals have a characteristic shiny appearance due to the presence of free delocalized electrons in the metallic structure. These electrons can absorb and re-emit light, giving metals their lustrous appearance.
Related to Metallic Bonding: In metallic bonding, metal atoms form a lattice surrounded by a “sea” of delocalized electrons that can reflect light, making the metal surface appear shiny.
2. Malleability (Ability to be Hammered into Thin Sheets)
Explanation: Metals are malleable, meaning they can be hammered or rolled into thin sheets without breaking.
Related to Metallic Bonding: The delocalized electrons in metallic bonding allow metal atoms to slide past each other without breaking the bond. This makes the metal flexible and able to change shape under pressure without fracturing.
Example: Gold and copper are often hammered into thin sheets for use in jewelry and electrical components.
3. Ductility (Ability to be Drawn into Thin Wires)
Explanation: Metals are ductile, meaning they can be drawn into thin wires.
Related to Metallic Bonding: Similar to malleability, delocalized electrons allow the metal atoms to shift positions without breaking the metallic bonds. This means the metal can be stretched into thin wires without losing its structural integrity.
- High Melting and Boiling Points
Metals are made up of a lattice of positive metal ions surrounded by a “sea” of free-moving electrons. The electrostatic attraction between these positively charged ions and the negatively charged electrons (metallic bonds) is very strong, requiring a lot of energy to break. - Heat Conductivity
Explanation: Metals are good conductors of heat, meaning they can transfer thermal energy efficiently.
Related to Metallic Bonding: The delocalized electrons in the metal lattice can move freely and carry thermal energy across the material. When heat is applied to one part of the metal, the delocalized electrons transfer the energy to other parts quickly. - Electrical Conductivity
Explanation: Metals are good conductors of electricity because they have free-moving delocalized electrons.
Related to Metallic Bonding: In metallic bonding, the electrons are not tightly bound to individual atoms. This allows them to move freely through the metal when an electric potential (voltage) is applied, facilitating the flow of electric current.
what is metallic bonding?
Metallic bonding is a type of chemical bonding that occurs between metal atoms. In metallic bonding, valence electrons (the outermost electrons of the metal atoms) are not tightly bound to individual atoms. Instead, these electrons become delocalized and move freely throughout the metal structure, forming a “sea of electrons.” This creates a strong bond that holds the metal atoms together in a regular lattice arrangement. In metallic bonding, metal atoms release their valence electrons into a shared “sea” or cloud of electrons. These electrons are free to move around and are not tied to any particular atom. This is what gives metals many of their characteristic properties.
what is oxidation
Oxidation is a chemical process in which an atom, ion, or molecule loses electrons. It is often accompanied by an increase in the oxidation state of the species involved. Oxidation is commonly part of a larger reaction called a redox reaction, in which one substance is oxidized (loses electrons) and another is reduced (gains electrons).
e.g. Fe → Fe²⁺ + 2e⁻
metal recycling
Mining and refining metals from the Earth.
Manufacturing products from refined metals.
Use of the products in various industries.
Disposal of products after they are no longer needed.
Recycling of metals, involving sorting, melting, and reprocessing.
Reprocessing into new or the same products, creating a closed loop of material reuse.
Resource Conservation: Recycling metals reduces the need for new mining, conserving natural resources and preventing habitat destruction.
Energy Savings: Recycling metals like aluminum and copper uses far less energy compared to extracting and refining new metals, reducing greenhouse gas emissions and the overall environmental footprint.
Waste Reduction: By recycling used products, fewer materials end up in landfills, reducing waste and contributing to environmental sustainability.
Economic Benefits: Recycling creates jobs, reduces costs in metal production, and can help keep prices more stable by reducing reliance on mined resources.
By embracing recycling, metals can be used sustainably, reducing the need for fresh mining, lowering energy consumption, and supporting a circular economy that minimizes waste and maximizes resource efficiency.
define alloy
mixture of two or more elements, where at least one of the elements is a metal. Alloys are created to improve the properties of the base metal, such as increasing strength, durability, resistance to corrosion, or enhancing other characteristics that make them more suitable for specific uses.
