Chem Exam 2024: The Chemical Awakens PDF

Chem Exam 2024: THe chemical awakens Coming in theatres December 10th Rated pg-100 Chemistry https://www.chemicalaid.com/tools/netionicequation.php?equation=Ni%28NO3%292+%2B+NaO H+%3D+Ni%28OH%292+%2B+NaNO3&hl=en Duration of exam: 1 hour 45 minutes Topics THE CHEMISTRY LAB, LABORATO...

Chem Exam 2024: THe chemical awakens Coming in theatres December 10th Rated pg-100 Chemistry https://www.chemicalaid.com/tools/netionicequation.php?equation=Ni%28NO3%292+%2B+NaO H+%3D+Ni%28OH%292+%2B+NaNO3&hl=en Duration of exam: 1 hour 45 minutes Topics THE CHEMISTRY LAB, LABORATORY EQUIPMENTS & SAFETY REGULATIONS PLANNING & DESIGNING EXPERIMENTS Identify and create questions and hypotheses that can be answered through scientific investigations. Develop appropriate experimental procedures for given questions/student generated questions. Analyse variables in scientific investigations soft Analyse evidence to explain observations, make inferences and predictions and develop the relationship between evidence and explanation. Present quantitative data resulting from scientific investigations: Use oral and written language to communicate findings, defend conclusions of scientific investigations and describe strengths and weaknesses of claims, arguments, and/or data tf dis mean?? CHEMICAL BONDING AND STRUCTURE Explain the formation of ionic and covalent bonds. ○ Ionic Bonds Forces of attraction caused from ions of opposite charges forming ionic bonds. Occurs between metals (like to lose electrons) and non-metals (like to take electrons) ○ Covalent Bonds Bonding with only non-metal atoms through the sharing of valence electrons Results in the formation of molecules Examples:  The binding force is due to the nuclei (plural of nucleus) of the atoms being attracted to the electrons being shared. Covalent bonds can be formed between two or more atoms of the same element. For example: ○ Fluorine molecule - Each fluorine atom contains unpaired electrons due to it only having 7 valence electrons (it is a halogen). Hence these lone pairs of electrons chose to form a bond that created the fluorine molecule.+ Lone pairs (non-bonding pairs) are electrons that are not involved in the bonding process. Fluorine has 3 lone pairs. ○ Oxygen Molecule - Each unbonded oxygen atom contains 2 unpaired electrons The 2 shared electron pairs represent 2 covalent bonds (double bond) Each oxygen atom in an oxygen molecule has 2 lone pairs ○ Nitrogen Molecule - Each nitrogen atom has 3 unpaired electrons Requires a triple bond to form 3 pairs of shared electrons representing a triple bond. Each nitrogen atom in a nitrogen molecule contains one lone pair of electrons.  The Molecular / Chemical Formula for covalent compounds shows the amount of each atom present in the molecule. Some covalent molecules formed between different atoms: ○ Water (H2O) ○ Ammonia (NH3) ○ Methane (CH4) ○ Carbon Dioxide (CO2) ○ Hydrogen Cyanide (HCN) Coordinate Covalent Bond Occurs when one of the atoms during the bonding process gives both electrons required to bond. The second atom contributes no electrons but instead they both bond by using the first atom’s lone pair. This forms a coordinate or dative covalent bond. ○ Metallic Bonding Bonds formed within metal atoms causing the valence electrons to leave forming cations. These electrons stay mobile and flow through the spaces between the positive ions.  Predict the likelihood of an atom forming an ionic or a covalent bond based on atomic structure. Write formulae to represent ions, molecules and formula units. Chemical and Ionic Equations ○ Chemical Equations A chemical equation is a shorthand representation of a chemical reaction. ○ Ionic Equations Ionic equations only show the atoms which actually take part in the reaction and as a result change state. Gyat gyat gytaatt- something suhavi does not have cookingwithkya Solubility & Insolubility Rules Compounds Containing Exceptions Solubility Rules Group 1 Elements (Li+, Na+, K+) None Nitrates (NO3-) None Ammonium (NH4+) None Perchlorates (ClO4-) None Hydrogen Carbonates None Acetates (C2H3O2-) & Ethanoates (CH3COO-) None Halogens such as: Chlorides (Cl-) Ag+, Pb2+, Cu2+, Hg22+ Bromides (Br-) Iodides (I-) Sulphates (SO42-) Group 2 Elements: Ca2+, Sr2+, Ba2+ Jerk metals: Ag+, Pb2+, Hg22+ Insolubility Rules Metal Hydroxides (OH-) Group 1 Elements Group 2 Elements: Ca2+ Sr2+, Ba2+ Ammonium (NH4+) Carbonates (CO32-) Group 1 Elements Phosphates (PO43-) Ammonium (NH4+) Sulphides (S2-) Group 1 Elements Group 2 Elements: Ca2+ Sr2+, Ba2+ Ammonium (NH4+) Metal Oxides Group 1 Elements Group 2 Elements: Ca2+ Sr2+, Ba2+ Monovalent Divalent Trivalent Cations Hydrogen - H+ Magnesium - Mg2+ Iron (iii) - Fe3+ Lithium - Li+ Calcium - Ca2+ Aluminium - Al3+ Sodium - Na+ Barium - Ba2+ Potassium - K+ Iron (ii) - Fe2+ Copper (i) - Cu+ Copper (ii) - Cu2+ Silver - Ag+ Zinc - Zn2+ Ammonium - NH4+ Tin (ii) - Sn2+ Lead (ii) - Pb2+ *Nickel (ii) - Ni2+ Anions Fluoride - F- Oxide - O2- Nitride - N3- Chloride - Cl- Sulphide - S2- Phosphate - PO23- Bromide - Br- Carbonate - CO32- Iodide - I- Sulphite - SO32- Hydride - H- Sulphate - SO42- Hydroxide - OH- Dichromate (vi) - Cr2O72- Nitrate - NO3- Manganate (vii) - MnO4- Hydrogen Sulphate - HSO4- Hydrogen Carbonate - HCO3-  Ethanoate - CH3COO- Chemistry Folder | Quizlet Describe ionic crystals, simple molecular crystals and giant molecular crystals; ○ The Structure and Properties of different elements and compounds depend on: The type of particles being formed - Crystals, gases, etc The force of attraction between the particles - Intermolecular and Intramolecular Intermolecular Forces ○ Forces between individual molecules ○ They are weak ○ Types: Van der Waal Forces A weak attraction between oppositely charged ends of a molecule. Hydrogen Bonds A weak bond forms between hydrogen and a more electronegative atom or group of another molecule  Intramolecular forces ○ Forces within the molecule (covalent bonds). ○ They hold the atoms together in the molecule. ○ They are strong ○ Structure & Properties of: Metals Ionic Crystals - Composed of an ionic lattice where cations and anions are held together in a regular repeating, 3D pattern by strong ionic bonds. Ionic Compounds  Simple Molecular Compounds / Crystals - Only a few atoms bonded together by strong covalent bonds. Arranged in a regular, 3D way to create a simple molecular lattice. The molecules within the lattice have weak intermolecular forces holding the small structure all together. Usually have low melting and boiling points. Referred to as simple covalent solids. Eg: Simple Covalent Compounds Particles can have the same element but different arrangements resulting in different forms. For example: Sulphur and Phosphorus Elements exhibit allotropy when they can exist in more than one form in the same physical state. Giant Molecular Compounds / Crystals - Compounds composed of non-metal atoms bonded by strong covalent bonds in a regular, 3D arrangement to form a giant molecular lattice. The covalent bonds exist between the atoms throughout the lattice causing it to be called a macromolecule as it is composed of millions of atoms.  Distinguish between ionic and molecular solids Relate structure of sodium chloride, diamond and graphite to their properties and uses Structure Sodium Chloride (NaCl) Diamond Graphite Properties 1) Hard, brittle crystalline 1) Each atom is bonded to 1) Each atom is arranged in solid. four others. flat 6 membered rings. 2) High melting and fusion 2) One of the hardest 2) Lubricating as the flat point. substances known. sheets slide across each 3) Conduct electricity well 3) Used in cutting and other. when molten or dissolved drilling. 3) Used in pencils and in water. 4) Does not conduct polishes. 4) Does not conduct electricity. 4) Conducts electricity. electricity in the solid state. 5) Most dissolve in water. 6) React with each other in water. Explain the term allotropy ○ Allotropes are different forms of the same element in the same physical state. Lil Summary Table Type of Occurs Involves… Particles formed Characteristics Properties of Structure of bonding between... of bond substances solid state formed Ionic Metals and Metals losing Cations Strong Hard, Giant ionic non-metals electrons and (positive) electrostatic crystalline Eg. Sodium non-metals attraction solids. chloride gaining electrons Anions between  Lil Summary Table (negative) oppositely High melting charged points. particles. Covalent Non-metal Sharing of Discrete Strong bond Usually liquids Usually simple atoms electrons. molecules between and gases. molecular molecule nucleus Eg. carbon May be single, Some can form and shared Solids formed dioxide, ice double or triple macromolecules. electrons have low bonds. (intramolecular). melting. Some are giant molecular Simple covalent Can be polar if Giant Eg. diamond, bonds only take the atoms molecules are graphite, silica one electron involved have hard with high from each atom. differing melting points. electronegativity. Coordinate covalent bonds Weak have both intermolecular electrons coming forces. from the same atom. Metallic Atoms of Atoms losing Positive ions and Electrostatic Hard, shiny, Giant metallic the same electrons which mobile electrons bond between electrical, metal become mobile positive ions and thermal electrons. conductors. TYPES OF REACTIONS AND BALANCING EQUATIONS Identify a reactant and a product in a chemical equation. State the law of conservation of matter ○ Matter and Atoms can neither be created nor destroyed in chemical reactions, they are only rearranged. Apply the law of conservation of matter when balancing chemical equations. identify and state how the four different state symbols are used: ○ Aqueous – (aq) ○ Solid – (s) ○ Liquid – (l) ○ Gas – (g) Identify the seven different types of chemical reactions. ○ Synthesis / Combination Reactions - Two or more substances are combined chemically to form a single product. ○ Decomposition Reactions - A single compound is broken down into two or more products. Occurs when a compound is unstable, heated or has an electric current go through it when the compound is in its liquid or dissolved state. ○ Single Displacement - When an element in its free state takes the place of another element in a compound. The more reactive element will always displace the less reactive element. Metals will displace other metals or displace hydrogen from its acid ie. Non-metals can displace other non-metals. ○ Double Displacement / Ionic Precipitation - This reaction occurs when two compounds which are in a solution exchange ions. One or more of the products must form a precipitate (salt). This prevents reversible reactions. ○ Neutralisation Reactions - A reaction between a base (alkali) and an acid. The acid is neutralised by the base and the products formed are salt and water. ○ Redox Reactions - This is a reaction in which one reactant is reduced and the other is oxidised. ○ Reversible Reactions - Occurs when the direction of a chemical change can be reversed easily. If the reaction is reversible a double arrow is used. ○ Combustion Reaction -  Oxygen combines with a compound to form carbon dioxide and water. These reactions are exothermic; give off heat. Write balanced ionic equations. PERIODIC TABLE AND PERIODICITY Explain the basis for the arrangement of elements in the periodic table  - Periodic table is arranged in terms of the atomic number (number of protons) - 18 groups (columns) and 7 periods (rows) - Elements in the same group have the same number of valence electrons. - The period number shows the number of shells an atom of an element has and what shell is the valence shell Periodic law: when the elements are arranged in order of increasing atomic number certain properties occur in a periodic manner. Periodic Trends The Periodic Table: Atomic Radius, Ionization Energy, and Electronegativity Atomic radius: measure of the size of an atom. The radius is measured from the nucleus to the valence shell. ○ Increase down a group New electron shell is added Inner shielding (the repel of the electrons from each other) ○ Slight decrease across a period Electrons are being added to the same shell Protons are being added therefore the nuclear pull/charge increases Greater attraction for outer electrons so there is a smaller radii (increase of effective nuclear charge)  Electronegativity: the ability of an atom to attract electrons to itself (the ability to gain electrons) ○ Measures the ability of an atom to attract electrons to form an anion so just the ability to gain electrons. ○ Decreases down a group and increases across a period ○ Inversely proportional to the atomic radii ○ Fluorine is the most electronegative element. So basically the atomic radius is small because it has the most amount of protons for its group with the same amount of shells so this means that the nuclear pull is strong. Since it has a strong pull it can gain electrons hence, it is the most electronegative element. Ionic Radii ○ Anions: have larger ionic radii because of electrons. Electron repulsion and weakened pull of the effective nuclear charge. ○ Cations: have smaller ionic radii because of less electrons. Electron repulsion and increased effective nuclear charge. Anions (Negatively Charged Ions): Larger Ionic Radii: Anions are atoms that have gained one or more electrons, resulting in a negative charge. Why They’re Larger: 1. More Electrons: When an atom gains electrons, it increases the total number of electrons surrounding the nucleus. 2. Electron Repulsion: The additional electrons lead to increased repulsion among the negatively charged electrons. This repulsion causes the electrons to spread out more, increasing the size of the electron cloud. 3. Weaker Effective Nuclear Charge: Although the number of protons (positive charge) in the nucleus remains the same, the additional negative charge from the extra electrons  reduces the pull the nucleus has on each electron. This weaker effective pull allows the electron cloud to expand, leading to a larger ionic radius. Cations (Positively Charged Ions): Smaller Ionic Radii: Cations are atoms that have lost one or more electrons, resulting in a positive charge. Why They’re Smaller: 1. Fewer Electrons: When an atom loses electrons, it decreases the number of electrons surrounding the nucleus. 2. Reduced Electron Repulsion: With fewer electrons, there is less repulsion among them. This allows the remaining electrons to be pulled closer to the nucleus. 3. Stronger Effective Nuclear Charge: The same number of protons is now pulling on fewer electrons. This means the effective nuclear charge (the net positive charge felt by the electrons) is stronger, pulling the remaining electrons closer to the nucleus and shrinking the ionic radius. Ionisation Energy: the energy needed to remove a valence electron from a neutral atom in the gaseous phase (1st ionisation energy) ○ Decreases down a group and increases across a period ○ 3 factors affect ionisation energy Distance of valence electrons from the nucleus. Increased distance means that the attraction between valence electrons and the nucleus decreases hence less energy is required to remove an electron. Effective nuclear charge The greater the nuclear charge/pull means the greater the attraction between valence electrons and the nucleus leading to more energy being needed to remove an electron. Shielding effect Outer electrons are shielded from the pull of the nucleus by inner electrons. The greater the shielding the less the amount of energy is required to remove an electron. What is Ionisation Energy? Ionisation Energy (IE) is the energy required to remove one electron from a neutral atom in the gaseous phase. The 1st Ionisation Energy refers to the energy needed to remove the first (outermost) valence electron. Trends in the Periodic Table: 1. Decreases down a group (moving from top to bottom): ○ As you go down a group, the ionisation energy decreases. 2. Increases across a period (moving from left to right): ○ As you go from left to right across a period, the ionisation energy increases. Factors Affecting Ionisation Energy: 1. Distance of Valence Electrons from the Nucleus: ○ The further the valence electron is from the nucleus, the easier it is to remove, leading to a lower ionisation energy. ○ This is why ionisation energy decreases as you move down a group; additional electron shells are added, increasing the distance between the nucleus and the outermost electrons. 2. Effective Nuclear Charge (Zₑff): ○ Effective nuclear charge refers to the net positive charge that an electron experiences from the nucleus. ○ Higher effective nuclear charge means that the nucleus exerts a stronger attractive force on the electrons, making it harder to remove a valence electron. This results in a higher ionisation energy. ○ Effective nuclear charge increases across a period because protons are added to the nucleus without a corresponding increase in inner-shell shielding, pulling the valence electrons closer. 3. Shielding Effect: ○ The shielding effect is the reduction in the effective nuclear charge felt by the outermost electrons due to the presence of inner-shell electrons. ○ Electrons in inner shells shield the outer electrons from the full charge of the nucleus, making it easier to remove a valence electron. ○ As you move down a group, more electron shells are added, increasing the shielding effect. This causes the outermost electrons to feel a weaker attraction from the nucleus, leading to lower ionisation energy. ○ However, across a period, the shielding effect remains relatively constant, so the effective nuclear charge increases and ionisation energy increases. Summary of Trends: Down a Group: Ionisation energy decreases due to increased distance and greater shielding effect. Across a Period: Ionisation energy increases due to increased effective nuclear charge with relatively constant shielding. Oxidizing Power: ability of an atom to cause another atom to lose its electrons. ○ Metals are reducing agents because they give their electrons to the non-metals ○ Non-metals are oxidizing agents because they gain the electrons that are lost from the metals. ○ Easier for metals to lose its valence electrons means it is stronger as a reducing agent and weaker as an oxidizing agent. So as u go go down the group for metals it's easier to lose electrons since the atom size is getting bigger  Metals are more reactive, stronger as reducing agents and weaker as oxidizing agents, so it is easier for them to lose their valence electrons. Explain trends in Group II ○ Beryllium, Magnesium, Calcium, Strontium, Barium, Radium Have similar properties, are very reactive when combined with other elements. ○ Reaction with oxygen Magnesium reacts slowly to form a coating of Magnesium Oxide (MgO) If ignited Mg burns with a blinding white flame to make a white, solid MgO. 2Mg (s) + O2(g) -- 2MgO(s) Calcium reacts readily to form a coating of Calcium Oxide (CaO) on exposure to air. If ignited, Ca burns with a brick red flame, making a white solid CaO. 2Ca (s) + O2(g) -- 2CaO(s) Barium reacts very readily to form a coating of Barium Oxide (BaO) on exposure with air. If ignited it burns with an apple green flame to make a white solid BaO. 2Ba (s) + O2(g) -- 2BaO(s) ○ Reaction with water Magnesium reacts very slowly with cold water to make Magnesium Hydroxide (Mg(OH)2 and Hydrogen (H2) Mg (s) + 2H2O(l) -- Mg(OH)2(s) + H2(g) Calcium reacts vigorously with cold water to make Calcium Hydroxide (Ca(OH)2) and Hydrogen gas (H2) Ca (s) + 2H2O(l) -- Ca(OH)2(s) + H2(g) Barium reacts very vigorously with cold water to make Barium Hydroxide (Ba(OH)2) and Hydrogen gas (H2) Ba (s) + 2H2O(l) -- Ba(OH)2(s) + H2(g) Reaction with Hydrochloric acid Magnesium reacts vigorously with HCL to make Magnesium Chloride (MgCl2) and Hydrogen gas (H2) Mg (s) + 2HCl(aq) -- MgCl2(aq) + H2(g) Calcium reacts very vigorously with HCL to make Calcium Chloride (CaCl2) and Hydrogen gas (H2) Ca(s) + 2HCl(aq) -- CaCl2(s) + H2(g) Barium reacts violently with HCL to make Barium Chloride (BaCl2) and hydrogen gas (H2) Ba (s) + 2HCl2(aq) -- BaCl2(s) + H2(g)  Metals react by preferably losing valence electrons to participate in bonding (ionization) If electrons are easier to remove, the element is generally more reactive, especially for metals. This is because: 1. Lower Ionisation Energy: Easier electron removal means lower ionisation energy, making metals more likely to lose electrons and form positive ions. This increases their reactivity. 2. Trends in Reactivity: ○ Down a Group: Reactivity increases as electrons are farther from the nucleus, making them easier to remove due to weaker attraction. ○ Across a Period: Reactivity decreases as ionisation energy increases from left to right, making electrons harder to remove. In summary, easier electron removal leads to higher reactivity, particularly for metals that lose electrons during reactions.  As the atomic radii increases, the valence electrons become further away from attractive pull of positive nucleus and the valence electrons become more shielded from nucleus The valence electrons are less attracted to nucleus, making it easier to be removed This ease of ionisation increases going down the group Explain trends in Group VII ○ They are poisonous ○ Non-metals ○ Exist as non-polar (no charge), diatomic molecules ○ Soluble in nonpolar solvents and slightly soluble in water. ○ Low melting and boiling points ○ Readily accept electrons into the valence shell to form a -1 anion. ○ They share electrons with other non-metallic atoms. ○ Electronegativity (ability to gain electrons) decreases down the group. Larger atomic radius and greater shielding down the group reduce the nucleus's attraction to bonding electrons, causing electronegativity to decrease from fluorine (most electronegative) to iodine (less electronegative). Chemical Reactivity: - React with most metals to form ionic compounds - 2Na(s)+Cl2(g) 2NaCl(s) - React with most non-metals to form covalent compounds - H2(g)+Cl2(g) 2HCl(g) - Reacts by gaining or sharing one electron  - Chlorine Water + Potassium Bromide Potassium Chloride + Bromine - Cl2(aq) + 2KBr(aq) 2KCl(aq) +Br2(aq) (red-brown) Identify trends in period 3 ○ They are solid at room temperature except mercury which is liquid. ○ They have high melting points and boiling points. ○ They have high densities. ○ They are good conductors of heat and electricity. ○ They are shiny in appearance. ○ They are malleable and ductile. Period 3 - General Properties ○ Non-metals ○ They are usually gases at room temperature however some are solids and bromine is liquid. ○ They have low melting and boiling points. ○ They have low densities. ○ They are poor conductors of heat and electricity. ○ In the solid state they are brittle and dull in appearance  Predict propertie s of unknown elements based on the position in periodic table