btec-national-applied-science-student-b1-unit-1
Principles and Applications of Science I
Stephen Coburn/123RF provided the photo credit.
UNIT 1: Principles and Applications of Science 1
This unit emphasizes the understanding of core science concepts for scientists and technicians. Chemists require knowledge of atoms and electronic structure to predict chemical reactions. Medical professionals need to understand cell structure and function for maintaining health and treating illnesses. Scientists in communication require wave knowledge.
Assessment
The unit is assessed through a 90-minute external paper worth 90 marks, set and marked by Pearson. The paper is divided into three sections:
Section A: Biology (Cell structure and function, Cell specialisation, Tissue structure and function)
Section B: Chemistry (Structure and bonding in applications of science, Production and uses of substances in relation to properties)
Section C: Physics (Working with waves, Waves in communication, Use of electromagnetic waves in communication)
The paper includes multiple-choice, calculations, short answer, and open-response questions, testing knowledge and understanding. Questions are contextualized to assess the application and synthesis of knowledge. The paper is available twice a year: January and May/June. Assessment practices throughout the chapter help prepare for the exam.
Assessment Outcomes (AO)
Unit 1 includes four Assessment Outcomes (AO):
AO1: Demonstrate knowledge of scientific facts, terms, definitions, and scientific formulae.
Command words: give, label, name, state
Marks range: 12 to 18
AO2: Demonstrate understanding of scientific concepts, procedures, processes, techniques, and their application.
Command words: calculate, compare, discuss, draw, explain, state, write
Marks range: 30 to 45
AO3: Analyze, interpret, and evaluate scientific information to make judgments and reach conclusions.
Command words: calculate, compare, comment complete, describe, discuss, explain, state
Marks range: 18 to 24
AO4: Make connections, use, and integrate different scientific concepts, procedures, processes, or techniques.
Command words: compare, comment, discuss, explain
Marks range: 9 to 12
Command Words and Definitions
Analyse: Identify relevant facts, demonstrate their links, and explain their importance.
Compare: Identify main factors of two or more items, point out similarities and differences, and state which are the best or most important.
Comment: Synthesize variables from data/information to form a judgment.
Define: State the meaning of something using clear and relevant facts.
Describe: Give a full account of all information, including all relevant details of any features.
Discuss: Write about the topic in detail, taking into account different ideas and opinions.
Evaluate: Bring all relevant information together and make a judgement on its success or importance, supported by gathered information.
Explain: Make an idea, situation, or problem clear by describing it in detail, including any relevant data or facts.
A Periodicity and Properties of Elements
A1: Structure and Bonding in Applications in Science
The Electronic Structure of Atoms
The nucleus of an atom contains positive protons and neutral neutrons. Energy shells surrounding the nucleus contain negative electrons. Protons and neutrons have a relative mass of 1, while electrons have a relative mass of almost 0. Lab technicians use knowledge of electronic structure to predict chemical behavior and reactions.
The electrons are in shells or energy levels surrounding the nucleus. Each shell can hold a maximum number of electrons, as shown in Table 1.1.
Table 1.1: Maximum Number of Electrons for Each Electron Shell
Shell 1: 2 electrons
Shell 2: 8 electrons
Shell 3: 18 electrons
Shell 4: 32 electrons
Shell 5: 50 electrons
A sodium atom with 11 electrons has an electron arrangement of 2, 8, 1. This is represented by a Bohr diagram. However, electrons within each shell have varying energy levels, divided into sub-shells called orbitals (s, p, d, and f). Orbitals have different energy states.
Orbitals: Regions where there is a 95% probability of locating an electron; each orbital can hold a maximum of two electrons.
The Aufbau principle states that electrons fill orbitals with the lowest available energy state closest to the nucleus before filling higher energy states, resulting in the most stable electron configuration. Electrons with the same charge repel each other, so they fill each orbital singly before pairing up.
Electron Structures
Electrons sit in orbitals within the shell.
The first shell holds two electrons in an s-type orbital.
The second shell has one s-type and three p-type orbitals.
The third shell has one s-type, three p-type, and five d-type orbitals.
Electrons fill the lowest energy level orbitals first.
Electrons occupy different orbitals when there are several of the same energy.
Electronic structure of nitrogen (7 electrons): 1s^2 2s^1 2p^1 2p^1 2p^1
Electronic structure of sodium (11 electrons): 1s^2 2s^2 2p^2 2p^2 2p^2 3s^1
Electron configuration: The distribution of electrons in an atom or molecule.
Spin: Electrons have two possible states, ‘spin up’ and ‘spin down’; each electron in an orbital will be in a different ‘spin state’.
Ionic Bonding
Noble gases have stable electronic configurations with full outer shells, making them unreactive. Other elements react to gain stable electronic configurations. Ionic bonding occurs when an atom loses electrons and donates them to another atom. The atom losing electrons becomes positively charged, and the atom gaining electrons becomes negatively charged.
For example, in sodium chloride (NaCl), sodium loses an electron to become Na^+, achieving the electron configuration of neon, while chlorine gains an electron to become Cl^-, achieving the electron configuration of argon. The electrostatic attraction between the oppositely charged ions holds them together.
Ionic bonding: Electrostatic attraction between two oppositely charged ions.
Ions containing more than one element can also form, such as the hydroxide ion (OH)− in sodium hydroxide (NaOH). The strength of the electrostatic force depends on the ionic charge and ionic radii of the ions. Higher ionic charge increases electrostatic force, while larger ionic radii weaken it due to charge distribution over a larger surface area.
Electrostatic attraction: The force experienced by oppositely charged particles; it holds the particles strongly together.
Giant ionic lattice: A regular arrangement of positive and negative ions, for example, in NaCl.
Covalent Bonding
Covalent bonding typically occurs between atoms of two non-metals, where an electron is shared between the atoms. These electrons come from the top energy level of the atoms.
A chlorine molecule (Cl_2) has a covalent bond where each chlorine atom shares one electron from its highest shell to achieve the electron configuration of argon, with a stable full outer shell.
Dative Covalent Bonding
In some covalent molecules, both sharing electrons come from one atom; this is a dative (coordinate) covalent bond. Oxygen atoms can form double bonds by sharing two pairs of electrons, while nitrogen molecules form triple bonds (N N).
An ammonium ion (NH_4^+) contains a dative bond. When ammonia reacts with hydrochloric acid, a hydrogen ion from the acid is transferred to the ammonia molecule. A lone pair of electrons on the nitrogen atom forms a dative covalent bond with the hydrogen ion.
Lone pair: A non-binding pair of electrons.
Single bonds have greater length than double bonds, and double bonds have greater length than triple bonds. Shorter bond length corresponds to stronger bonds. For example:
Single bond (C-C): length 154 pm, energy 347 kJ/mol
Double bond (C=C): length 134 pm, energy 612 kJ/mol
Triple bond (C≡C): length 120 pm, energy 820 kJ/mo. l
Covalent Bonding in Organic Molecules
Carbon makes four covalent bonds, forming many organic compounds. Methane (CH_4) has each carbon atom bonding covalently with four hydrogen atoms, giving carbon the stable electron structure of neon and hydrogen of helium. Methane has a tetrahedral structure because the bonds are separated as much as possible, with each bond angle being 109.5^o.
Organic compound: A compound that contains one or more carbons in a carbon chain.
Organic compounds with three or more carbons in a chain cannot be linear due to the tetrahedral structure around each central carbon.
Metallic Bonding
Metals have giant structures of atoms held together by metallic bonds in a regular lattice. Metallic bonding occurs because electrons in the highest energy level of a metal atom are delocalized, moving freely through the metal in a ‘sea’ of electrons. This gives the metal nuclei a positive charge, attracted to the negative charge on the delocalized electrons.
Delocalised electron: Electrons that are free to move; present in metals and not associated with a single atom or covalent bond.
Forces in metallic bonding are strong but not as strong as in covalent or ionic bonding. The electronegativity of two atoms determines the type of bond formed between them.
Electronegativity: The tendency of an atom to attract a bonding pair of electrons.
Atoms with similar electronegativities form covalent bonds with a strong electrostatic attraction between the nuclei and the shared pair(s) of electrons. For example, chlorine molecules (Cl_2) are non-polar as both atoms have the same electronegativity.
In most covalent compounds, the bonding is polar covalent, where shared electrons are attracted more to one nucleus, giving it a slight negative charge, and the other atom a slight positive charge. As the difference in electronegativity increases, the bond becomes more polar.
Non-polar molecule: A molecule where electrons are distributed evenly throughout.
Polar molecule: A molecule with a partial positive charge in one part and a partial negative charge in another due to uneven electron distribution.
The electronegativities of common elements include:
Fluorine: 3.98
Oxygen: 3.44
Nitrogen: 3.04
Carbon: 2.55
Chlorine: 3.16
Hydrogen: 2.20
Lithium: 0.98
Sodium: 0.82
Intermolecular Forces
Intermolecular forces affect chemical substance behavior. Laboratory technicians must understand their presence and effects on chemical substances.
London Dispersion Forces
These are weak forces (less than 1% of a covalent bond) present between non-polar covalent molecules, also called temporary dipole-induced dipole forces. When electron distribution in a molecule becomes asymmetrical, temporary dipoles form and induce dipoles in nearby molecules, causing attraction. Larger molecules have more electrons and can form bigger dipoles, leading to a stronger attraction.
More electrons → more movement → bigger dipoles → stronger attraction
Intermolecular forces: Attraction or repulsion between neighboring molecules.
Dipole: Separation of charges within a covalent molecule.
London dispersion forces are the only forces existing between noble gases and non-polar molecules.
Dipole-Dipole Forces
These are permanent forces between polar molecules, slightly stronger than London dispersion forces but still weak (about 1% the strength of a covalent bond). Polar molecules have permanent positive and negative ends that attract each other. Molecules with dipole-dipole forces include hydrogen chloride (HCl) and iodine monochloride (ICl).
Van der Waals forces: All intermolecular attractions are van der Waals forces.
Hydrogen Bonding
This is the strongest intermolecular force (about 10% the strength of a covalent bond) and occurs when hydrogen is directly bonded to fluorine, oxygen, or nitrogen. The large electronegativity difference creates very polar bonds and permanent dipoles. A hydrogen bond forms between the positive end of one molecule and the lone pair of electrons of another, strongly drawing the hydrogen to the nitrogen, oxygen or fluorine atom. Hydrogen bonding in water causes its unusual properties, such as higher boiling point and being a good solvent.
Quantities Used in Chemical Reactions
Balancing Equations
All chemical reactions can be written as a balanced equation using the chemical formulae for the reactants and products involved in the reaction following these steps:
Write the equation as a word equation including all the reactants and all the products.
Write out the formulae for each substance in the reaction. Note that gaseous elements (except those in group 0) like hydrogen and oxygen are diatomic (molecules with two atoms) so they must be written as H2 and O2. Metal elements and the noble gases are monatomic (one atom)
Write out the number of each element on both sides.
Make the number of each atom equal on each side. Remember that you cannot change the formula of the compounds. To increase the number of atoms of a particular element, you must place a number in front of the compound it is in. This will affect the number of atoms of all the other elements in the compound.
Check that there is the same number of atoms of each element on both sides.
Example: Ethanol + oxygen → carbon dioxide +water
C2H5OH + O2 2CO2 + H_2O becomes
C2H5OH + 3O2 2CO2 + 3H_2O
Write a balanced equation for the following reaction: butanol (C4H9OH) + water → carbon dioxide + water.
Now write a balanced equation for: magnesium carbonate + hydrochloric acid → magnesium chloride + water + carbon dioxide.
Moles, Molar Masses and Molarities
Chemical equations allow calculation of reactant masses needed for specific product masses. Chemists use moles (mol), where one mole equals 6.023 × 10^{23} particles (Avogadro’s constant).
A mole is the amount of a substance which has the same number of particles as there are atoms in 12 g of carbon-12. So one mole of carbon dioxide has the same number of particles as one mole of gold.
The molar mass of a substance is equal to the mass of one mole of a substance. Convert masses into moles and moles into masses using the following equation:
Mass (g) = molar mass × number of moles ∴ Number of moles = mass/M_r
Mole: A unit of substance equivalent to the number of atoms in 12 g of carbon-12; 1 mole of a compound has a mass equal to its relative atomic mass expressed in grams.
Molar mass: The mass of one mole of a substance.
The relative atomic mass (A_r) of an element on the periodic table tells you how much mass there is in one mole of the element. The relative atomic mass is the average mass of an atom of an element compared to one twelfth of the mass of an atom of carbon–12. The relative atomic mass of hydrogen is 1.0. The relative atomic mass of oxygen is 16.0.
The relative formula mass is the sum of all the relative atomic masses of all the atoms in the empirical formula (simplest formula) of a compound (M_r).
The relative formula mass of water, H_2O, is (1 × 2) + 16 = 18.
Relative atomic and formula masses do not have any units as they are only relative to carbon–12.
Divide the mass of each element present in the compound by its molar mass to get its molar ratio
Divide the answer for each element by the smallest molar ratio calculated. This gives you a ratio of 1:x for each element present.
If the answers are not all whole numbers, multiply them all by the same number to get whole numbers. e.g. if the ratio is 1:1.5:3 then multiplying all the numbers by 2 will give you an answer with all whole numbers 2:3:6
Types of Formula
Empirical Formula: Shows the ratio between elements in a chemical compound, useful for giant structures like sodium chloride.
Molecular Formulae: Used for simple molecules. To work out the molecular formula you need to know the empirical formula and the relative molecular mass.
Working it out.
A compound has the empirical formula CH2. It has an empirical formula mass of 12 + (1 × 2) It has a relative molecular mass of 42. To work out its molecular formula you first divide its relative molecular mass by the empirical mass. 42/14 = 3 You write out the formula multiplying each part of the CH2 unit by 3. This gives C3H6. This is the molecular formula.
Reacting quantities
Titrations - a chemist has to use solutions of a known concentration. These are called standard solutions. They have been prepared and tested to ensure they are of the specific concentration needed
The number of moles of solute in a given volume of solvent tells you how concentrated the solution is
1 mole of solute is dissolved in 1 cubic decimetre of solution, its concentration is written as: 1 mol dm−3. This can be written as 1M for short. This is the molarity of the solution. 36.5 g of HCl in 1 dm^3 of solution has a concentration of 1 mol dm^{-3} or 1M or 36.5 g dm^3.
Titration: A method of volumetric analysis used to calculate the concentration of a solution.
Solution: A liquid mixture where a solute is dissolved in a solvent.
Standard solution: A solution of known concentration used in volumetric analysis.
Solute: The substance dissolved in a solvent to form a solution.
Solvent: A liquid that dissolves another substance.
Number of moles (N) = molarity (C) × volume of solution (V) (dm^3) : N = CV The volume is given in cm^3 so this needs to be converted into dm^3 by dividing by 1000. (Remember 1 dm^3 = 1000 cm^3)
Using a Chemical Equation to Calculate the Quantities of Reactants and Products
Used to calculate the quantities of reactants and products.
Calcium chloride can be produced by reacting calcium carbonate with hydrochloric acid. This equation includes state symbols:
CaCO3 (s) + 2HCl (aq) CaCl2 (aq) + CO2 (g) + H2O (l)
This is an example of stoichiometry.
Stiochiometry: Involves using the relationships between the reactants and the products in a chemical reaction to work out how much product will be produced from given amounts of reactants.
Ar (H) = 1, Ar (C) =12, Ar (0) = 16, Ar (Cl) = 35.5, A_r (Ca) = 40
One mole of CaCO3 produces one mole of CaCl2, A 1:1 ratio.
40 +12 + (3 × 16) g = 100 g of CaCO3 produces 40 + (35.5 × 2) g = 111 g of CaCl2.
As one mole of CaCO3 produces one mole of CaCl2 then:
100 g CaCO3 produces 111 g CaCl2.
The equation for the reaction follows:
2H2 +O2 2H_2O
Add up the relative atomic masses for each substance. Remember there will be two lots of water.
2 × 16 g = 32 g of O2 produces 2 × (2 × 1) +16 g = 36 g of H2O
32 g O2 produces 36 g H2O
So 10 g O_2 produces ___ 36 32 × 10 g of H_2O
Percentage Yields
The theoretical mass is the amount of product you can produce in a reaction. In most reactions it is unlikely that the total amount of product possible is made
Theoretical mass: The expected amount of product from a reaction calculated from the balanced equation.
Percentage yield: The actual amount of yield worked out as a percentage of the theoretical yield.
Percentage yield = \frac{Actual mass}{Theoritical mass} \times 100
actual number of moles × 100
Reversible reaction: A reaction where the reactants react to form products and the products simultaneously react to re form the reactants.
A2 Production and Uses of Substances in Relation to Properties
The Periodic Table
The periodic table shows all the chemical elements arranged in order of increasing atomic number. It can be used to predict element behavior, and physical or chemical properties. A laboratory technician needs to be very familiar with the periodic table, and organized into groups (vertical columns) and periods (horizontal rows). Elements in the same group have similar chemical properties. The atomic number increases from left to right across a period because each successive element has one more proton than the one before.
Characteristics of each period:
Period 1
Contains hydrogen and helium, both gases. Electrons fill the 1s orbital. Helium is unreactive. Hydrogen loses or gains an electron and can behave chemically as both a group 1 and a group 7 element. Hydrogen can form compounds with most elements and is the most abundant chemical element in the universe.
Period 2
Contains eight elements: lithium, beryllium, boron, carbon, nitrogen, oxygen, fluorine, and neon. The outer electrons fill the 2s and 2p orbitals. N, O, and F can form diatomic molecules. Neon is a noble gas. Carbon is a giant molecular structure.
Period 3
Contains eight elements: sodium, magnesium, aluminum, silicon, phosphorus, sulfur, chlorine, and argon. The outer electrons fill the 3s and 3p orbitals.
Period 4
Contains 18 elements, from potassium to krypton. The first row of the transition elements is in this period. The outer electrons fill the 4s, 4p, and 3d orbitals.
Groups – s block, p block, d block
The periodic table is organized by element blocks, named for the orbital the highest energy electrons are in.
Groups 1 and 2 are in the s block.
Groups 3 to 7 and group 0 make up the p block. This block contains all the non-metals except for hydrogen and helium.
The transition metals are in the d block.
The highest energy electron in carbon is in a p orbital and therefore carbon is a p block element.
Physical Properties of Elements
Atomic Radius
The atomic radius decreases across the period (left to right) due to increased nuclear charge attracting electrons closer to the nucleus. Atomic radii increases down a group as the extra electrons are added to additional shells, and nuclear charge is shielded more.
Ionic Radius
Ionic radius trends down a group follow a similar pattern to the atomic radius trend; extra electrons are added to extra shells giving a larger size. Cations have a smaller radius than their corresponding atom. They are isoelectronic and are all arranged with nuclear charge. Anions have a larger radius than the corresponding atom because there is more repulsion between the extra electrons.
Isoelectronic: Having the same numbers of electrons.
Cations: Ions with a positive charge.
Anions: Ions with a negative charge.
Electronegativity
Electronegativity increases across a period and decreases down a group, making fluorine the most electronegative element. This depends on the number of protons in the nucleus, the distance from the nucleus of the bonding pair of electrons and the amount of shielding from inner electrons.
First Ionisation Energy
It is the minimum energy needed for one mole of the outermost electrons to be removed from one mole of atoms in a gaseous stated:
For example, the equation shows potassium losing one electron to become a positive ion.
K(g) K^+(g) + e^- .
It shows Periodicity, increasing across the period and decreasing down the group. The increased positive nuclear charge more strongly attracts the outer electron, leading to a subsequent higher ionization energy
Periodicity: The repeating pattern seen by the elements in the periodic table.
First ionisation energy: The energy needed for one mole of electrons to be removed from one gaseous atoms.
Electron Affinity
It is an atom’s ability to gain an electron and become a negative ion. It is the change in energy (kJ mol–1) of a neutral gaseous atom when an electron is added to the atom to form a negative ion.
It is affected by nuclear charge, distance from nucleus, and shielding by inner electrons. This is because they are all factors to the attraction involved when it comes to an electron to the nucleus of an atom in the atom state. Fluorine is different as it is a very samll atom - so there is repulsion from these.
O(g) + e− O−(g). O(g) + e− → O−(g).
Electron affinity: The charge in energy when one mole of a gaseous atom gains one mole of electrons to form a negative ion.
Type of bonding in the element
The electronegativity of elements can be used to predict the type of bonding in a compound. Bonding is a spectrum from ionic to covalent
Elements in periods 1 to 3 and their Electronegativities
H 2.1
Li 1.0 Be 1.5 B 2.0 C 2.5 N 3.0 O 3.5 F 4.0
Na 0.9 Mg 1.2 Al 1.5 S 1.8 P 2.1 S 2.5 Cl 3.0
Polarised or Non-Polarised!
Ionic bonds can also show polarity depending on :
Either ion is highly charged
The cation is relatively small
The anion is relatively large
Trends: melting point and boiling point
The elements in the periodic table also show periodicity for melting and boiling points. Melting and boiling points depend on the strength of the forces between the atoms in an element.
Period 2
As you go across groups 1 to 3, metals have increasing nuclear charge because they have increasing number of protons and increasing number of delocalised electrons and so have stronger metallic bonding. This means the melting and boiling points increase as you go across the metals in the period.
Carbon, has giant covalent bonding forming a giant lattice structure with each atom bonding to 4 other carbon atoms so its melting and boiling points are very high because it has strong covalent bonds that need a large amount of energy to break. The non-metals in groups 5 to 7 have small separate molecules and so have low melting points, there are only weak van der Waals forces that need to be overcome.
Period 3
It follows the simple trend of period 2 with a few small exceptions. Sulfur, in group 6, has a higher melting and boiling point than the rest of the non-metals. This is because of the different size of the molecules of each of these elements. Phosphorus exists as P4, sulfur as S8, chlorine as Cl_2 molecules and argon as Ar atoms. The strength of the van der Waals forces increases as the size of the molecule increases. Therefore, because sulfur has the biggest molecule, it has the strongest van der Waals forces and so the highest melting and boiling points.
Physical Properties of Metals: Electrical Conductivity, Thermal Conductivity, Malleability, Ductility
Metallic bonding allows for electrical conductivity through a solid or liquid metal, and the delocalised electrons carry the electric charge. Delocalised electrons in metals also absorb heat energy and then are transferred through the metal making them good thermal conductors. The structure of metals can also explains why they can be malleable or ductile- as the atoms are able to roll over each over and can move to new positions without breaking the metallic bonds.
Chemical Properties of Elements
The reactions between oxygen and metals are very important due to its influenace and easily of reaction. Iron reacts very easily with oxygen and forms rust so it is often painted to protect it from oxygen in the air.
The following are the reaction equations and a few notes on the nature of each one;
Lithium: 4Li (s) + O2(g) → 2Li2O(s)
Rapid, burns with red flame. The metal oxide produced forms a basic solution when dissolved in waterSodium: 4Na (s) + O2(g) → 2Na2O(s); 2Na (s) + O2(g) → Na2O2(s)
Very vigorous, burns with orange flame. Metal oxide produced that form basic solution when dissolved in waterBeryllium and magnesium: 2Be (s) + O2(g) → 2BeO(s)
Needs heat to react as do group 1 elements, also very vigorous reactions.Aluminium: 4Al (s) + 3O2(g) → 2Al2O3(s)
Vigorous at first, rapidly forms a water insoluble coating of Al203. This layer prevents the aluminium below from corroding and so makes aluminium an extremely useful material.Carbon Silicon: C(s) + O2(g) → CO2(g) ; 2C(g) + O2(g) → 2CO (g)
Forms slightly acidic oxides, and silicon shows reaction with heat.Nitrogen and Phosphorus: The equation and amount of oxides released depends on temperature
Forms a range of oxides with different oxidation states, A high temperature is need for these reactions to take place.Oxygen and Sulfur: O + O2 → O3 03 S + O2 → SO2
In ozone layer with two oxides forming.Most halides react: Not usually formed by direct reaction.
Unstable oxides form.Neon Argon: No reaction as they are noble gasses.
Metals and O2 reaction
Oxides formed change and are described to go from left to right with the product changes also changing in character to become from solids to gasses on the right.
Reactions with oxygen is faster more violently depending on the position on the Periodic table.
For example:
Lithium, sodium and potassium are stored under oil to prevent contact with air due to this or they are stored in sealed glass tubes to ensure that no air or oxygen is present.
Their oxides contain the simple ion O^{2– }. Sodium and potassium can also form the peroxide M2O2 containing the molecular ion to which are unstable to the O2 2–. The covalent bond between the two negative oxygen ions in O_2^{2-}
is weak.Lead and tin also production oxides with formula MO and MO_2
Lead/Tin oxides are often brittle
The Reactivity series will be less reactive for group d and some become even corrosion residents, as a unreactive outer oxide layer prevents any more of the metal's reaction.
Other Examples of Oxidation/Reactions with oxygen
4M + 3 O2 2M2O_3
Beryllium often forms beryllium oxidation
Beryllium/ Aluminium behvave as unreactive metals
Zinc d- block metals oxides made with oxygen are often brittle
Reactions with water
Group 1 are called alkali metals because they produce a basic solution. Metals react when using reaction series.
2M(s) +2H2O(l) 2M^+(aq) +2OH^–(aq) +H2(g)
Group 2 magnesium only reacts with steam, while the metals below magnesium will react increasingly easily with water and do produce hydroxides in the equation.
Beryllium does not react with water.
There is a similar reaction with group aluminium - the outer aluiminium oxide layer prevents it
M(s) +2H2O(l) M(OH)2(aq) +H_2(g)
Group 4, 5 and 6 metals do not react with water, transition metals react slowly with water and some do not react at all.
Reactions with Dilute Acids
Metals above copper in the reactivity series can react with dilute acids to form metal salts and hydrogen:
Mg +2HCl → MgCl2 + H2.
It reacts with dilute sulfuric acid to give magnesium sulfate and hydrogen:
Mg +H2SO4 → MgSO4 + H2.
Sodium reacts with hydrochloric acid to form sodium chloride and hydrogen:
2Na + 2HCl → 2NaCl + H2.
The Reactivity Series
Reacticity of a metal is its ability to form a complete outer Shell to lose an electron. From which the more reactive is the most likely to be in found as a compound form.
The most reactive metals are in group 1, and decreasing across the table, and increase in reactivity further down group 1 from Francium. However Francium is excluded from as its to radioactive
The d-block elements are less reactive and the order of reactivity is group 1, group 2, group 4, transition metals.
Reduction, Oxidation and Oxidation states
The term redox refers to the transfer of electrons that occurs during chemical reactions. When atoms lose electrons, it is called oxidation and when they gain electrons it's called reduction. (Use the mnemonic 'OIL RIG' to remember this is 'Oxidation is Loss', 'Reduction is Gain' ). This process of electron transfer allows the reaction and oxidation and reduction to occur simultaneously.
Mg Mg^{2+} +2e^–
(Magnesuim has been oxidized)
½ O_2 + 2e^– O^{2−}
(Oxygen has been reduced
There cant be a hafl equation for the