Foundations In Chemistry Flashcards
State the 3 sub atomic particles their properties
Protons
relative charge +1
relative mass 1
Neutrons
Relative charge 0
Relative mass 1
Electrons
Relative charge -1
Relative mass 1/2000 (negligible)
Define isotope
Atoms of the same element which have different numbers of neutrons
Define relative atomic mass
The weighted mean mass of an atom of an element compared to 1/12 the mass of a carbon 12 atom
Define relative isotopic mass
The mass of an atom of an isotopes compared to 1/12 of a carbon 12 atom
Define ion
An ion is an atom that has a charge due to gaining or losing an electron(s)
What is the mass number and atomic number
Mass number (the biggest number-MASSive) is equal to the number of protons+neutrons in the nucleus
Atomic number is equal to the number of protons in the nucleus (also equal to the number of protons in ATOMS NOT IONS)
To find the number of neutrons you subtract the atomic number from the mass number
What is the calculation for relative atomic mass?
RAM= (abundance 1 x isotopic mass 1) + (abundance 2 x isotopic mass 2) / 100
Define shells, sub shells and orbitals
Shells- a shell is a group of atomic orbitals with the same principle quantum number (shell number). The shell closest to the nucleus has a number of 1 and has the lowest energy.
Sub shells- a group of the same type of atomic orbitals within a shell (s, p, d, f)
Orbitals- a region in atom around the nucleus which can hold up to two electrons with opposite spins.
Describe the shape of s and p orbitals
S orbitals- spherical shape
P orbitals- dumbbell shape
Describe the filling of atomic orbitals
1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^10 4p^6
What is the electron configuration for chromium (24 e-) and copper (29 e-)?
Why are they like this?
Chromium: 1s^2 2s^2 2p^6 3s^2 3p^6 4s^1 3d^5
Copper: 1s^2 2s^2 2p^6 3s^2 3p^6 4s^1 3d^10
This is because they are more stable this way
How do u draw the filling of atomic orbitals?
Boxes with arrows (you don’t want stinky people next to you)
In what order do electrons fill orbitals?
Electrons fill from the lowest energy level orbitals first up to the highest energy level orbitals last
Why is X classed as an (s/p/d/f) block element?
Because it’s highest energy level electron occupies an (s/p/d/f) orbital
Define transition elements
A d Block Element which forms an ion with an incomplete d sub shell
What do transition elements lose first when forming positive ions?
4s electrons
Define empirical formulae
The simplest whole number ratio of elements in a compound
How do you calculate empirical formulae
Calculating Empirical Formula:
1. Determine the mass of each element in the compound.
2. Convert the masses to moles using the relative atomic masses.
3. Divide the moles of each element by the smallest number of moles to get the ratio.
• Example: For a compound containing 6.0 g of carbon and 8.0 g of hydrogen, the moles of carbon and hydrogen are calculated, and the ratio is used to find the empirical formula.
Define molecular formulae
The actual number of atoms of each element in a molecule
How do you calculate molecular formulae
The molecular formulae can be determined if the relative atomic mass of the compound is known:
• Formula: molecular formulae = Empirical formula x n
• Where n is the ratio of the molar mass of the compound to the empirical formula mass.
What are the rules for metal ions, non metal ions and transition metal ions
Metal ions: +ve ions usually ending in -ium
Non-metal ions: -ve ions usually ending in -ide, (-ate or -ite which mean containing oxygen)
Transition metals: +ve ions of variable charge
What is an ionic compound and how do you write them
Consist of oppositely charged ions held together by electrostatic forces
-Identify the charges on the ions.
-Balance the charges to form a neutral compound via the crossover method
Name the different compound ions
OH (-)
NO3 (-)
MnO4 (-)
CO3 (2-)
SO4 (2-)
SO3 (2-)
Cr2O7 (2-)
PO4 (3-)
How do you write chemical equations
-Write the equation in words
-Put in all the state symbols and chemical formulae
-Balance the equation
Define acid
An acid is a proton or H+ donor. They release H+ ions in an aqueous solution.
Explain the difference between a strong and weak acid
Strong Acids: Completely dissociate in aqueous solutions, releasing all its acidic hydrogen atoms
Weak Acids: Only partially dissociate in aqueous solutions, releasing some but not all of its acidic hydrogen atoms
Define base and list some types
A base is a proton or H+ acceptor.
Types:
Metal oxide
Metal hydroxide
Metal carbonate
Ammonia solution
Define an alkali
An alkali is a base that dissolves in water and releases OH- ions in aqueous solutions
Define a salt
A salt is produced when the H+ ion of an acid is replaced by a metal or NH4+
It is produced from the neutralisation of an acid and a base
Explain the reactions to make a salt
Acid + ammonia = ammonium salt
Acid + metal oxide = salt + water
Acid + metal hydroxide = salt + water
Acid + metal = salt + hydrogen
Acid + metal carbonate = salt + water + carbon dioxide
Explain the solubility of salts
-All sodium, potassium, and ammonium salts are soluble.
-Most chloride salts are soluble, except for silver chloride and lead(II) chloride.
-Most sulfate salts are soluble, except for barium sulfate and calcium sulfate.
Explain the properties of a salt
Salts are generally crystalline and have high melting and boiling points due to the strong ionic bonds.
Salts like sodium chloride can conduct electricity when molten or dissolved in water because the ions are free to move.
Explain what a standard solution is and how to produce one
A standard solution is a solution of known concentration
To prepare it:
-Accurately weigh the solid (mass by difference method)
-Dissolve the solid in a beaker using a small amount of distilled water
-Carefully transfer the solution into the volumetric flask. Ensure you rinse the
beaker with distilled water and add this to the volumetric flask.
-Carefully fill the flask with distilled water until the bottom of the meniscus
lines up exactly with the graduation mark
-Invert the flask several times (with the stopper fitted) to mix the solution.
Explain how to carry out an acid-base titration
-25.0cm3 of the solution A was transferred into a conical flask using a pipette
-Indicator was added to the solution in the conical flask
-Solution B of unknown concentration was added to the burette and the initial
burette reading was recorded to the nearest 0.05cm3
-The solution B was slowly added to the conical flask until the indicator
changed colour at the end point.
-The final burette reading was recorded to the nearest 0.05cm3
-The titration was then repeated until two concordant results were achieved
-record the data in a table, containing the initial volume, final volume and titre
What is an acid base titration used for
To find the concentration of a solution of unknown concentration from a solution of a known concentration
How do you calculate the unknown concentration
Concentration 1 x volume 1 = concentration 2 x volume 2 (given the units are the same)
This is because the moles are the same because there is an equal number of moles of acid to moles of the base which is why it becomes neutralised
Explain what moles are
A mole is an amount of a substance which is equivalent to the number of atoms in 12g of carbon 12 or 6.02x10^23 particles (atoms, molecules etc)
How do you calculate moles
Moles=mass/mr
Explain molar mass
The mass in grams per mole of a substance
How do you calculate number of particles
Number of particles = moles x Avagadros constant
How do you calculate moles of a substance using stoichiometric ratios
-Write a Balanced Chemical Equation. Ensure the equation is balanced before starting calculations.
-Identify the Known and Unknown
-Use the Mole Ratio from the balanced equation, determine the stoichiometric relationship between the known and unknown substances
-Perform Calculations. Convert given quantities (mass, volume, concentration) into moles if necessary.
-Use the mole ratio to calculate the number of moles of the unknown substance.
-Convert moles back to required units (mass, volume, etc.) if needed.
How do you calculate moles in a solution
Moles = concentration x volume (divide by 1000 if one of the units is in cm3 or by 1000000 if both are in cm3)
How do you calculate moles of a gas under standard condition and non standard conditions
-pV=nRT (non standard conditions)
Pressure in pascals
Volume in meters cubed
n in moles
R constant = 8.314
Temperature in K (to convert from degrees Celsius +273)
-n=V/24
n in moles
Volume in dm cubed
-n=V/24000
n in moles
V in cm cubed
Explain percentage yield and how to calculate it
Percentage yield measures the efficiency of a chemical reaction by comparing the actual yield of a product to the theoretical yield
Percentage yield = actual yield/theoretical yield (x100)
Explain why you may have a low yield
-Incomplete reactions (not all reactants react).
-Side reactions (producing unwanted by-products).
-Loss of product during purification or transfer.
-Reversible reactions not going to completion.
Explain atom economy and how to calculate it
Atom economy measures the proportion of reactants that are converted into useful (desired) products. It evaluates the sustainability and efficiency of a chemical reaction.
Atom economy = Mr of desired products/Sum of Mr of all reactants (x100)
Explain why atom economy is relevant
-Atom economy indicates how much waste a reaction produces.
-Reactions with high atom economy are more sustainable and environmentally friendly
-Higher atom economy reduces waste and improves efficiency, making industrial processes more sustainable
Explain how to maximise atom economy
-Use reactions where most/all reactants become the desired product.
-Reactions with fewer by-products have higher atom economies.
Explain the difference between atom economy and percentage yield
Percentage yield focuses on the actual efficiency of producing the product compared to theoretical amounts whereas atom economy focuses on the proportion of reactants that end up in the desired product, regardless of losses or inefficiencies
Explain what oxidation and reduction are
Oxidation is a loss of electrons and reduction is the gain of electrons
Explain what an oxidising agent is and what a reducing agent is
Oxidising agents accept electrons and get reduced
Reducing agents donate electrons and get oxidised
What are the general rules for oxidation numbers
-Uncombined elements + HONClBrIF = 0
-ions = charge of the ion
-Hydrogen = +1 unless in a metal hydride where its -1
-Oxygen = -2 unless in peroxides where its -1 or bonded to fluorine where its +2
How do you find the oxidation state of an element in a compound
Start with the most electronegative and progress down until you finish with all your known oxidation states. Then assign an appropriate oxidation state, accounting for the charge of the compound
How do you identify what is being oxidised or reduced
Oxidation: increase in oxidation number
Reduction: decrease in oxidation number
How do you write a half equation
-Identify species being oxidised or reduced
-Balance atoms (except H and O)
-Balance oxygen using H₂O
-Balance hydrogen using H⁺ ions
-Balance charge using electrons (e⁻)
How do you combine half equations
-Ensure the number of electrons in oxidation and reduction half-equations are equal
-Add the two half-equations together and cancel electrons
-simplify
What is the term for when a substance is oxidised and reduced simultaneously
Disproportionation
What are the exceptions for oxidation numbers
H in metal hydride (-1)
O in peroxides (-1)
O bonded to F (+2)
Define ionic bonding
The electrostatic attraction between oppositely charged ions formed by the transfer of electrons from a metal to a non-metal
Describe the formation of ionic bonds
-Metal atoms lose electrons → form positive ions (cations).
-Non-metal atoms gain electrons → form negative ions (anions).
The electrostatic forces between oppositely charged ions are strong and extend throughout the lattice.
Describe the properties of ionic compounds
High Melting and Boiling Points
-Strong electrostatic forces require a lot of energy to break.
Conductivity
-Solid state: Does not conduct (ions are in a fixed position).
-Molten or dissolved: Conducts (ions are free to move).
Solubility
-Generally soluble in polar solvents like water (polar molecules surround and separate ions).
-Insoluble in non-polar solvents.
Describe covalent bonding
The strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms.
Describe dative covalent bonding
The strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms where one atom donates both electrons in the bond.
Describe electronegativity and its trend on the periodic table
The ability of an atom to attract a bonding pair of electrons in a covalent bond. The greater its ability to attract, the greater its electronegativity.
-Increases across a period (more protons, smaller atomic radius).
-Decreases down a group (larger atomic radius, more shielding).
It increases on the periodic table from the bottom left corner up to the right corner (Francium is the least electronegative and Fluorine is the most electronegative).
Describe polar and non polar bonds
-Non polar bonds have equal electronegativity due to equal distribution of electrons.
-Polar bonds have unequal electronegativity due to an uneven distribution of electrons. This creates a permanent dipole across the bond.
Describe symetrical molecules and use CCl4 as an example
Symmetrical molecules are non polar because the dipoles cancel out. Example: CCl4 is a symmetrical molecule. Each C–Cl bond is polar because there is a difference in electronegativity. However, because the CCl4 molecule is symmetrical, the dipoles cancel out. Therefore the CCl4 molecule is non polar as there is no net dipole across the molecule.
Describe asymmetrical molecules using NH3 as an example
Unsymmetrical molecules, which contain one or more polar bonds, are polar because the dipoles do not cancel out. Example: NH3 is an asymmetrical molecule with polar N–H bonds. Each N-H bond is polar because N is more electronegative than H. Since the NH3 molecule is asymmetrical, the dipoles do not cancel. Therefore the NH3 molecule is polar.
Define intermolecular forces and state the different types
Intermolecular forces are attractive forces between molecules. Intermolecular forces are much weaker than ionic or covalent bonds. They are only found in covalent structures. Intermolecular
forces can be:
-hydrogen bonds
-permanent dipole-dipole interactions
-induced dipole-dipole interactions
Describe induced dipole dipole interactions
These are very weak intermolecular forces found between all molecules.
At any moment, there may be an uneven distribution of electrons in a molecule. This causes a temporary dipole to be present, causing an induced dipole in neighbouring molecules. The δ+ of a dipole in one molecule attracts the δ - of a dipole in a neighbouring molecule to produce a London dispersion force.
Describe permanent dipole dipole interactions and use HCL as an example
These are weak attractive forces between polar molecules. Example: hydrogen chloride, HCl. The H–Cl bond is permanently polar because Cl is more electronegative than H, and the molecule is asymmetrical
so the dipoles don’t cancel.
The Hδ+ of one HCl molecule attracts the Clδ- of a neighbouring HCl molecule to produce a permanent dipole dipole force of attraction between the molecules.
Permanent dipole-dipole attractions are stronger than induced dipole-dipole attractions.
Describe hydrogen bonding
A hydrogen bond is a strong dipole-dipole attraction between molecules containing O-H, N-H or F-H
bonds (FOHN). A hydrogen bond exists between a Hδ+ atom in one molecule and a lone pair on a highly electronegative atom on another molecule. Hydrogen bonds are the strongest type of intermolecular forces of attraction.
Explain the anomalous properties of water due to hydrogen bonding
Ice (solid H2O) is less dense than liquid water. In the solid (ice) H2O molecules are held further apart by the hydrogen bonds. This gives the ice an open lattice structure, and so it’s less dense.
The H2O molecules are fixed in their positions (cannot move around) in
the 3-dimensional ice crystal lattice by the hydrogen bonds.
Water has relatively high melting and boiling points. Water has a higher than expected melting point and boiling point. Hydrogen bonds are relatively strong so more energy is needed to break/overcome them.
Describe the structure and properties of giant ionic lattices.
Structure: Regular arrangement of oppositely charged ions (e.g., Na⁺ and Cl⁻).
Properties:
-High melting/boiling points due to strong electrostatic forces between ions.
-Conducts electricity when molten or in solution (ions are free to move).
-Soluble in polar solvents (e.g., water) as they can separate ions.
Example: NaCl (sodium chloride).
Describe the structure and properties of giant covalent lattices.
Structure: Atoms bonded in a continuous network by strong covalent bonds (e.g., carbon atoms in diamond).
Properties:
-High melting/boiling points due to strong covalent bonds.
-Does not conduct electricity (except graphite where delocalized electrons are present).
-Insoluble in water due to strong covalent bonds.
Examples:
-Diamond: Hard, tetrahedral structure, does not conduct electricity.
-Graphite: Layers of hexagonal rings, conducts electricity due to delocalized electrons.
Describe the structure and properties of simple molecular structures.
Structure: Small molecules held together by weak intermolecular forces (e.g., CO₂, H₂O).
Properties:
-Low melting/boiling points due to weak intermolecular forces.
-Does not conduct electricity (no free electrons or ions).
-Soluble in non-polar solvents (e.g., CO₂ in organic solvents).
Examples: CO₂, O₂, H₂O.
Explain metallic bonding and its properties
The strong electrostatic attraction
of a lattice of positive metal ions to
a ‘sea’ of delocalised electrons.
Properties:
-Electrical Conductivity: The free-moving electrons carry charge, allowing metals to conduct electricity.
-Thermal Conductivity: The electrons also transfer heat energy efficiently.
-Malleability and Ductility: The layers of metal ions can slide past each other without breaking the bond, making metals easy to shape (malleable) or stretch (ductile).
-High Melting and Boiling Points: The strong attraction between the metal ions and the delocalised electrons requires a lot of energy to break, giving metals high melting and boiling points.
-Luster (Shiny Appearance): The free electrons reflect light, giving metals their shiny look.
Explain electron pair repulsion theory
-Electron pairs repel each other to get as far apart as possible
-Lone pairs repel more strongly than bonding pairs
Describe all the shapes of molecules, their bond angles and any bonded/lone pairs
-Linear: 2 bonded pairs, 180 degrees
-Trigonal planar: 3 bonded pairs, 120 degrees
-Tetrahedral: 4 bonded pairs, 109.5 degrees
-Octahedral: 6 bonded pairs, 90 degrees
-Non linear: 2 bonded pairs, 2 lone pairs, 104.5
-Trigonal Pyramidal: 3 bonded pairs, 1 lone pair, 107 degrees
For each lone pair, subtract 2.5 degrees from the bond angle
