G9 Term 2 Science ETA Reviewer Chemistry PDF

TERM 2 CHEMISTRY ETA Reviewer Organized by Grade 9L Cycle 1:​ Lesson 1: Atomic Structure Atoms - Atoms are the smallest units of matter that retain the ch...

TERM 2 CHEMISTRY ETA Reviewer Organized by Grade 9L Cycle 1:​ Lesson 1: Atomic Structure Atoms - Atoms are the smallest units of matter that retain the characteristics of an element, composed of three subatomic particles: protons (positive charge), neutrons (neutral), and electrons (negative charge). They cannot be divided further by chemical or physical means. Protons Positively charged particles, and are found in the nucleus of an atom and determine the element's identity. Electrons Negatively charged particles, and are spread out around the nucleus of an atom and are responsible for creating compounds through chemical bonding. Neutrons They have no charge and do not change the properties of an atom. They serve as a buffer between protons in the nucleus, helping to stabilize the atom. Atomic Number The number of protons in the nucleus of an atom. Mass Number Total number of protons and neutrons in an atom's nucleus. To find: ​ Protons: The atomic number is the total number of protons. ​ Electrons: The number of electrons is equal to the number of protons. ​ Neutrons: Subtract the atomic number from the mass number to find the number of neutrons. Atomic No. of No. of No. of Mass ELEMENT Number protons neutrons electrons Number Titanium (Ti) 22 22 26 22 48 Yttrium (Y) 39 50 39 Bismuth (Bi) 83 83 209  Cycle 1: Lesson 2: The Periodic Table of Elements Elements - Pure substances made of one type of atom, defined by the number of protons in the nucleus. Can be classified as metals, nonmetals, and metalloids​ Periods - Horizontal rows in the Periodic Table, indicating the number of electron shells.​ Groups or families - Vertical columns in the Periodic Table, containing elements with similar properties and the same number of valence electrons. The periodic table is composed of: ​ Seven (7) horizontal rows (periods) ​ Eighteen (18) vertical rows (groups/families) Groups in the Periodic Table Description G1: Alkali Metals Highly reactive and naturally exist as compounds with other elements. Have low densities and melting points. G2: Alkaline Earth Metals Highly reactive, solid, harder than alkali metals, and good conductors of heat and electricity. Denser with higher boiling and melting points than Group 1 elements. G3 - 12: Transition Metals Hard, shiny solids with high thermal and electrical conductivity and high melting points. Often form multiple charged states. G13: Boron Group (Earth Metals) Elements with three valence electrons, including metals and metalloids, some of which are semiconductors. G14: Carbon Group (Tetrels) Elements with four valence electrons, including nonmetals, metalloids, and metals; essential for organic compounds. G15: Nitrogen Group (Pnictogens) Elements with five valence electrons, including nonmetals and metalloids; commonly form nitrogen-based compounds. G16: Oxygen Group (Chalcogens) Elements with six valence electrons, including oxygen and sulfur, essential for life and forming various compounds. G17: Halogens Highly reactive nonmetals that react violently, especially with metals. Poor conductors of heat and electricity and have low boiling and melting points. G18: Noble Gases Colorless, odorless, and nonflammable under standard conditions, most stable elements. Inert gases with eight valence electrons, making them stable and nonreactive. ​ ​ Reading the periodic table  ​ ​ ​ ​ ​ Metals, Nonmetals, Metalloids Cycle 1: Lesson 3: Periodic Trends Atomic Radius -​ the side of an atom or radius in the distance from the nucleus to the boundary of the surrounding electron cloud. -​ decreases as you move from left to right across a period. -​ increases as you move down from a group. Nuclear attraction -​ As the number of protons increases, the force of attraction between the nucleus and electrons increases, shrinking the atom. Ionic Radius -​ This is the size of the ion as the atoms gain or lose electrons. -​ When atoms lose electrons and form Cations which are positively charged ions, they become smaller. -​ When atoms gain electrons and form Anions (negatively charged ions), they become larger. Ionization energy -​ It's the energy that's required to remove an electron from a gaseous atom to form a gaseous ion. This energy represents how strongly an atom holds onto its outermost electrons. -​ Increases across a period (left to right) as atoms have more protons and a greater nuclear charge. -​ Decreases down a group (top to bottom) as electrons are farther from the nucleus and more shielded by inner electrons. Electronegativity -​ the ability of an atom to attract electrons towards itself when forming a chemical bond. -​ Increases across a period (left to right) as atoms attract electrons more strongly.  -​ Decreases down a group (top to bottom) as atomic radius increases. Summary: Cycle 2: Lesson 4: Quantum Mechanical Model of Atom Atomic Theory (Democritus) -​ All matter is composed of tiny indivisible particles called atomos, he came up with the atomic theory. Dalton’s Atomic Theory (John Dalton) -​ All elements are made up of tiny particles called “atoms”. -​ All atoms of a particular element are identical. -​ The atoms of one element are different from another element. -​ Atoms of an element can combine with atoms of another element to form compounds. -​ Atoms are indivisible, they can’t be created nor destroyed in chemical reactions. Thomson’s Plum Pudding Model (JJ Thomson) -​ Discovered electrons, negative particles within atoms. -​ Proposed the atom is a "plum pudding" with electrons embedded in a positively charged "pudding." Rutherford’s Gold Foil Experiment (Ernest Rutherford) -​ Positive charge is not like a pudding but it's concentrated in the nucleus. -​ Concluded that atoms have a dense, positively charged nucleus with electrons surrounding it. Bohr’s Orbit Model (Niels Bohr) -​ electrons in an atom occupy certain orbits with quantized energies. -​ Atoms have allowable energy states -​ Electrons revolve in stable orbitals around a nucleus -​ Atoms have certain allowable energy states Heisenberg’s Uncertainty Principle (Werner Heisenberg) -​ Possible to assign a fixed path for electrons. -​ It's impossible to know exactly where an electron is at any moment. -​ There is uncertainty in the position and motion of the electron. Quantum Mechanical Model of An Atom -​ It is the map of the probably location of the electrons -​ It helps us predict the orbital’s possible location, shape, and energy of the location In order to get the location of an electron in an atom, we need: ​ Energy Level: The fixed amount of energy an electron can absorb, with 7 main energy levels. ​ Sublevel (Subshell): Defines the shape and type of orbitals. ​ Orbitals: Regions where electrons are calculated to be present. Outermost Energy Level -​ Valences electrons are electrons found in the outermost energy level of an atom -​ These are the electrons that can be gained or lost during a chemical reaction. Principle Sublevels (Types of Number of Orbitals Total Number of Orbitals Quantum Number Orbitals) Present Related to Sublevel Related to Principal Energy (n) Level (n2) 1 s 1 1 2 s 1 p 3 4 3 s 1 p 3 9 d 5 4 s 1 p 3 16 d 5 f 7  Cycle 2 Lesson 5: Electronic Configuration Aufbau Principle -​ Electrons that fill the lowest energy orbitals first before moving to higher energy ones. -​ The order of orbitals that are filled follows this sequence: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p, and so on. Pauli Exclusion Principle - Only two electrons may occupy an orbital, and they must have opposite spins Hund's Rule - States that each orbital must receive one electron before any orbital can receive a second filling of electron. Determining the Electron Configuration: 1.​ Determine the element’s atomic number 2.​ Know the kind of sublevel per energy level 3.​ Add subscript for each letter designation to indicate the number of electrons per orbital 4.​ Follow the Aufbau and Pauli Exclusion Principle Cycle 4 Lesson 6: Ionic and Covalent Bonding Valence Electrons -​ electrons in the last energy level of an atom -​ can be gained or lost during a chemical reaction -​ number of valence electrons = group number Lewis Dot Structure -​ represents the number of valence electrons in an element -​ Valence Electrons = dots = no side can have more than 2 dots = each side gets one dot before doubling up Chemical Bonding ​ How atoms combine to form compounds ​ The type of chemical bond present in a compound dictates its physical and chemical properties ​ Electrostatic forces of attraction holds atoms together Electrostatic Forces - Forces that either attract or repel based on charges.  Electrostatic Forces Attract or Repel Attract ​ Repel Types of Chemical Bonding Ionic Bond Transfer of electrons Covalent Bond Sharing of electrons Metallic Bond Overlap of orbitals/delocalized electrons Ionic Bonds -​ exists between two oppositely charged ions (cations and anions) -​ an element loses electrons while another element gains it -​ exists between a metal and nonmetal -​ metal loses valence electron → nonmetal gains the electron to achieve octet -​ metal loses electrons = cation is formed -​ nonmetal gains electrons = anion is formed -​ elements w/ low electronegativities + elements w/ high electronegativities = ionic compounds Covalent Bonds -​ atoms share electrons -​ happens between nonmetals -​ exists when the differences in electronegativities is not enough for electron transfer -​ Single Covalent Bond = sharing of a single pair of electrons = valence electron not involved in bonding - Lone Pairs -​ Double Covalent Bond = sharing of 2 pairs of electrons -​ Triple Covalent Bond = sharing of 3 pairs of electrons  Electronegativity Type of Bond Example Differences (△EN) △EN < 0.4 Nonpolar Covalent Bond Cl2 1.7 > △EN > 0.4 Polar Covalent Bond HCl △EN > 1.7 Ionic Bond NaCl Nonpolar Covalent Bonding -​ electrons are shared equally between atoms -​ generally insoluble in water Polar Covalent Bonding -​ electrons are unequally shared between atoms -​ tend to be soluble in water Metallic Bonds -​ held together by the strong attraction of positive cations and negative electrons -​ malleable and ductile -​ good conductors of electric current Cycle 4 Lesson 7.1 & 7.2: Naming and Formula Writing of Ionic and Covalent Compounds Ionic Compounds (Metal + NonMetal) Writing Ionic Formulas: 1.​ Identify cation (positive ion) and anion (negative ion). 2.​ Balance charges by crisscrossing (use subscripts to balance the charges). Example: Na⁺ + Cl⁻ → NaCl, Mg²⁺ + Cl⁻ → MgCl₂. Naming Ionic Compounds: 1.​ Name cation first (no prefix). 2.​ Name anion second with “-ide.” 3.​ For transition metals, use Roman numerals for charge. Example: NaCl = Sodium chloride, FeCl₂ = Iron (II) chloride. FeCl3 = Iron (III) chloride Covalent Compounds (Nonmetal + Nonmetal) Writing Covalent Compounds: 1.​ Use prefixes for the number of atoms: mono-, di-, tri-, etc. Example: CO₂ = Carbon dioxide, N₂O₅ = Dinitrogen pentoxide. Naming Covalent Compounds: 1.​ Name the first element with a prefix (if more than one atom). 2.​ Name the second element with the prefix and change ending to “-ide.” Example: CO₂ = Carbon dioxide, N₂O₅ = Dinitrogen pentoxide. Prefixes for Covalent Compounds: 1 mono 2 di 3 tri 4 tetra 5 penta 6 hexa 7 hepta 8 octa 9 nona 10 deca Properties of Ionic and Covalent Compounds Ionic Covalent Metal + Nonmetal Nonmetal + Nonmetal Transfer of electrons Sharing of electrons Crystalline solids (made of ions) Gases, liquids, or solids (made of molecules) High melting/boiling points Low melting/boiling points Conducts electricity in water Does not conduct electricity Many soluble in water but not in nonpolar Many soluble in nonpolar liquids but not in liquids water or other polar liquids Uses charges in naming Uses prefixes in naming Cycle 5 Lesson 8: Evidences of Chemical Reaction and Balancing Chemical Equations Law of Conservation of Mass -​ Mass cannot be created or destroyed in a chemical reaction, but is conserved -​ The total mass of reactants is equal to the total mass of products. -​ Atoms are rearranged, but they are neither created nor destroyed. Chemical Equation -​ a statement used by chemists to represent chemical reactions. -​ shows the reactants and products of a reaction and the direction in which the reaction progresses. Word Equation -​ Uses words to describe a chemical reaction -​ “Aluminum and Chlorine reacts to form Aluminum Chloride Skeleton Equation -​ uses symbols to describe a chemical reaction -​ ex. 2Al(s) + 3Cl2(g) → 2AlCl3(s) -​ symbols: (s) = solid, (l) = liquid, (g) = gaseous, (aq) = aqueous Example Skeleton Equation: ​ ​ Step 1: Write the Skeleton Equation​ ​ Step 2: Count the number of atoms of the elements in the reactant side.​ ​ Step 3: Count the number of atoms of the elements in the product side.​ ​ Step 4: Check if the number of atoms are the same in both sides of the chemical equation.​ ​ Step 5: If the number of atoms aren’t the same, add coefficients to make it the same.​ ​ Note: DO NOT TOUCH ANY SUBSCRIPTS.​ ​ Cycle 5 Lesson 9: Types of Chemical Reactions and Factors Affecting Reactions Combination or Synthesis -​ A type of chemical reaction in which two or more simple substances combine to form a more complex substance. -​ A + B → AB Decomposition -​ a compound breaks down two or more simpler compounds or elements -​ AB → A + B Single Replacement -​ A + BC → AC + B -​ an element replaces another element in a compound Double Displacement -​ 2 compounds exchange cations to produce 2 different compounds -​ AB + CD → AD + BC Combustion -​ happens when a hydrocarbon reacts with oxygen, producing energy in the form or heat and light -​ Complete Combustion = happens when sufficient amount of oxygen reacts in the reaction  = products of it are carbon dioxide(H2O) and water(H2O) = CxHy + O2 → CO2 + H2O -​ Incomplete Combustion = happens when sufficient amount of oxygen reacts in the reaction = products of it are carbon(soot)(C)/ and or carbon monoxide(CO) and water(H2O) = CxHy + O2 → C + CO + H2O Rate of Chemical Reaction -​ how fast or slow the reactants are converted into products -​ related to the activation energy of a chemical reaction Activation Energy -​ minimum amount of energy needed to start a chemical reaction -​ higher activation = slower chemical reaction -​ lower activation = faster chemical reaction 1.​ Increase the surface area of the reactants -​ more surface area exposed = faster reaction 2.​ A. Increase the temperature of solids and liquids -​ higher temperature = faster particle movement = faster particle interaction -​ lower temperature = slower particle movement B. Decrease the temperature of gases -​ cooler temperature, slower particle movement 3.​ Increase the amount or concentration of the reactants -​ increased amount of reactants allows for more particles to react 4.​ Addition of catalyst -​ Catalyst = a material that increases the rate of reaction by lowering the activation energy Evidence of Chemical Reactions: 1.​ Formation of gas (bubbles) 2.​ Formation of precipitate (solid) 3.​ Change in pH 4.​ Color change 5.​ Temperature change 6.​ Production of light 1 4Al + 3O2 → 2Al2O3 combination 2 2KClO3 → 2KCl + 3O2 decomposition 3 2Al + Fe2O3 → Al2O3 + 2Fe Single replacement or displacement 4 2NaOH + H2SO4 → Na2SO4 + 2H2O Double replacement or displacement 5 [C2H6 + 3.5O2 → 2CO2 + 3H2O] x 2 2C2H6 + 7O2 → 4CO2 + 6H2O Complete combustion