9-CONCHEM Past Paper PDF

Q1 9-CONCHEM Q2 9-CONCHEM Q2 - Chemistry Ma’am Mary Edd Anne Certeza | 2nd Quarter | Owned by Ysabelle Ng, Lindsay Inocando, Stephen Villarin, Stephanie Trinidad UNIT IV: Order Among Elements Chemical Symbols Development of Periodic Table Arrangement of Elements in the Periodic...

Q1 9-CONCHEM Q2 9-CONCHEM Q2 - Chemistry Ma’am Mary Edd Anne Certeza | 2nd Quarter | Owned by Ysabelle Ng, Lindsay Inocando, Stephen Villarin, Stephanie Trinidad UNIT IV: Order Among Elements Chemical Symbols Development of Periodic Table Arrangement of Elements in the Periodic Table Electron Configuration Group No. & Valence electrons Quantum Numbers Trends in the Periodic Table The Chemical Behavior of Elements UNIT V: Stoichiometry Chemical formula Mole concept ○ Mole conversion ○ Molar Mass ○ Percentage Composition Empirical and Molecular Formula Chemical equations ○ Balancing Chemical Equations Types of Chemical Reactions Quantitative information from balanced equations Limiting reactants and theoretical Yields UNIT VI: Chemical Bond Chemical bond, Lewis symbol, octet rule Ionic, covalent, metallic bond Bond polarity and electronegativity Geometry of molecules Molecular polarity Intermolecular forces of attraction UNIT IV: Order Among Elements 4.1-3 The Periodic Table (10/16/24) Periodic Table Tabular arrangement of chemical elements organized according to increasing atomic number The Historical Development of the Periodic Table 1817: Johann Wolfgang Döbereiner’s Law of Triads Arranged elements into groups of three based on their similar properties, called triads ○ Examples: Lithium, sodium, potassium ○ The average of the atomic masses of the first and third elements is the atomic weight of the 2nd element Ex: (Atomic mass of Lithium + Atomic mass of Potassium) / 2 = Atomic mass of Sodium 1865: John Newlands’s Law of Octaves Every eighth element has similar properties when the elements are arranged based on their increasing atomic masses This is only true for the elements until calcium (Ca) 1869: Dmitri Mendeleev’s Periodic Table Created the first periodic table  ○ Elements were arranged by increasing atomic mass ○ Left gaps or spaces for elements that were yet to be discovered Mendeleev is known as the Father of the Periodic Table 1913: Henry Moseley’s Periodic Table Modern Periodic Table Law ○ Chemical and physical properties of elements tend to vary by increasing atomic number, not atomic mass Parts & Features of the Periodic Table 1. Periods Horizontal rows in the periodic table There are 7 periods 2. Groups/Family Vertical columns in the periodic table Total of 18 families ○ Consists of Family A (1A-8A, also called representative elements) and Family B (transition metals) Family A: (idk if u need to memorize this but yeah) ○ 1A: Alkali Metals ○ 2A: Alkaline Earth Metals ○ 3A: Boron Group ○ 4A: Carbon Group ○ 5A: Nitrogen Group/Pnictogens ○ 6A: Oxygen Group/Chalcogens ○ 7A: Halogens ○ 8A: Noble Gases 3. Types of Chemical Elements 4. Oxidation State 5. Valence Electrons 6. Blocks 4.4 Electron Configuration (10/16/24) The arrangement or distribution of electrons in atomic orbitals ○ Recall: Atomic orbital - where an electron can most likely be found ○ Most stable state: Ground state Recall: Ground state happens when all energy levels are full ○ Unstable state: Excited state (when electrons jump to another energy level) Recall: Energy level - distance of electrons from the nucleus How to write electron configuration: 1. Energy level - (lowest/closest to nucleus) 1, 2, 3, 4, 5, 6, 7 (highest/ farthest from nucleus) Note: Theoretically, there are no limitations to the energy level but, an atom rarely have electrons beyond the 7th energy level. 2. Type of orbital/sublevel - s, p, d, f s (sharp) - spherical ○ maximum of 2 electrons p (principal) - dumbbell ○ maximum of 6 electrons d (diffuse) - clover ○ maximum of 10 electrons f (fundamental) - flower-shaped ○ maximum of 14 electrons 3. Number of electrons in the orbital - written in superscript 1. Aufbau principle Electrons first fill the subshells of the lowest available energy, then fill subshells of higher energy unless the atom is in an excited state. 2. Pauli exclusion principle A maximum of two electrons may occupy an orbital if they have opposite spins. ○ This is because electrons are negatively charged, and electrostatic force will cause them to repel. So in an orbital, one electron should spin in a clockwise direction and the other will spin in a counterclockwise direction to prevent repulsion. 3. Hund’s Rule Every orbital is singly occupied before any orbital gets doubly occupied. Shorthand method (using noble gas configuration) Only used for elements in periods 2 and above. This method does not apply to elements in period 1 (hydrogen and helium) We use the noble gas prior to the period of the element ○ For example, we have to use helium for lithium’s electron configuration as that is the noble gas before period 2, the period lithium is in ○ Example: Sodium’s electron configuration: 1s22s22p63s1 The underlined configuration is Neon’s electron configuration. Therefore, Sodium’s electron configuration (shorthand method): [Ne] 3s1 ○ Example: Fluorine’s electron configuration is 1s22s22p5 The underlined configuration is Helium’s electron configuration. Therefore, Fluorine’s electron configuration (shorthand method): [He] 2s22p5 Following order to fill the orbitals: 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6, 5s2, 4d10, 5p6, 6s2, 4f14, 5d10, 6p6, 7s2, 5f14, 6d10, 7p6 4.5 Valence Electrons and Blocks (10/18/24) Valence Electrons The electrons in the last energy level or outermost shell. ○ Octet Rule Elements must attain 8 valence electrons to be stable. Exceptions are hydrogen and helium, instead, they follow the duplet rule (2 electrons instead of 8) Valence electrons can tell us the family number for representative elements (family A) Family 1A 2A 3A 4A 5A 6A 7A 8A Orbital s1 s2 p1 p2 p3 p4 p5 p6 + no. of e- Valence 1 2 3 4 5 6 7 8 Electro ns Charge +1 +2 +3 ±4 -3 -2 -1 0 s Explanation for the Table: For an element to be stable, it must follow the octet rule that states: An element must attain 8 valence electrons to be stable. For example, in Family 1A, the valence electron is 1 but to be stable, it needs to be 8. To attain this stability, the element must gain 7 electrons or lose 1 electron. Losing 1 electron is easier than gaining 7 electrons, hence it has a +1 charge.  ○ Note: If an atom loses an electron, it gains a positive charge. If an atom gains an electron, it gains a negative charge Cation: positive ion Anion: (-) ion How does losing an electron attain stability? (for Valence Electrons 1-4): ○ Take, for example, Family 1A with Valence electron 1 and a +1 charge. We know that a +1 charge means to lose 1 electron; losing an electron allows the element to attain the electron configuration of the previous noble gas which follows the octet rule, thus attaining stability. ○ Example: Li has an electron configuration of 1s22s1, belongs to Family 1A, and has a valence electron of 1. → We know that to attain the stability of an element with v.e. 1, it must have a +1 charge (lose 1 electron). → If we remove 1 electron from 1s22s1, we get 1s2. 1s2 is the electron configuration of Helium, which is a noble gas. Things to remember: Valence electrons tell the group number of elements in Family A only ○ For example: Since Lithium is in group 1A, it has 1 valence electron Periods tell the last energy level ○ For example: Since Lithium is in period 2, its last energy level is 2 Blocks in the Periodic Table: The periodic table is divided into four main blocks—s, p, d, and f—based on the type of atomic orbital that the element’s valence electrons occupy. s-block Groups: 1 and 2 Valence electrons: Occupy the s-orbital p-block Groups: 13 to 18  Valence electrons: Occupy the p-orbital d-block (Transition Metals) Groups: 3 to 12 Valence electrons: Occupy the d-orbital f-block (Lanthanides & Actinides) Placement: Often shown separately below the main table. Valence electrons: Occupy the f-orbital. Example: Element Electron configuration Group Valence Period Energy Electrons level Magnesium 1s22s22p63s2 2A 2 3 3 (Mg) Hydrogen (H) 1s1 1 1 1 1 If given the Group and Period: Element Electron configuration Group Valence Period Energy Electrons level 8A (given) 2 (given) We know that energy level tells us the period and vice versa. Thus, if we have 2 as our given period, our energy level must be 2. Element Electron configuration Group Valence Period Energy Electrons level 8A (given) 2 (given) 2 We know that the valence electrons identify the group and vice versa. Thus, if we have 8A as our given group, our valence electrons must be 8. Elements found in period 2, family 8A have a p orbital. This gives us 2p and although we have 8 valence electrons, we cannot write it as 2p8, but we know that p has a maximum of 6 electrons giving us 2p6. We can now form the electron configuration until we reach 2p6. This gives us: 1s22s22p6. Adding up the electrons, we get 10 which is also our atomic number. Thus, our element that belongs to family 8A period 2 is Neon (Ne). Element Electron configuration Group Valence Period Energy  Electrons level Neon (Ne) 1s22s22p6 8A (given) 8 2 (given) 2 Getting the Electron Configuration, Group #, V.E., Period #, and Energy Level of Family B Elements Element Electron configuration Group Valence Period Energy Electrons level Yttrium (39) Step 1: Write the electron configuration Element Electron configuration Group Valence Period Energy Electrons level Yttrium (39) 1s22s22p63s23p64s 3B 2 3d104p65s24d1 Step 2: To get the Group #, find the no. of electrons in the d-orbital and add 2. In this case, the number of electrons in the d-orbital is 1; 1+2 = 3. Thus, it’s in Family 3B Element Electron configuration Group Valence Period Energy Electrons level Yttrium (39) 1s22s22p63s23p64s 3B 3 2 3d104p65s24d1 Step 3: To get the valence electrons, add the no. of electrons in the s-orbital and d-orbital. 2+ 1 = 3. Thus, there are 3 valence electrons Element Electron configuration Group Valence Period Energy Electrons level Yttrium (39) 1s22s22p63s23p64s 3B 3 5 5 2 3d104p65s24d1 Step 4: Getting the period number is similar to getting the period number of Family A elements. The outermost shell is 5, thus it’s the period number. The period number determines the energy level. *There are exceptions. Such as Copper (Cu), you would need to add the electrons in s-orbital and d-orbital according to the EXPECTED electron configuration ([Ar] 4s23d9), not the ACTUAL electron configuration ([Ar] 4s13d10). 9+2=11, but since the d-orbital can handle a maximum of 10 electrons, you subtract 10 from 11 and get 1. Hence, Copper is found in Family 1B. 4.6 Electromagnetic Radiation and Quantum Numbers (10/21/24) Electromagnetic spectrum The full range of electromagnetic radiation according to frequency and wavelength Includes (from higher energy to lower): ○ Gamma-ray ○ X-ray ○ Ultraviolet ○ Visible Light ○ Infrared ○ Microwave ○ Radio waves ○ Ultralow frequency All light is part of the EM spectrum. Amplitude: The brightness/intensity of light depends on the amplitude of light wave Frequency: How fast a wave completes a cycle of upward and downward motion per second Measured in hertz (Hz) Wavelength: Distance between two successive peaks (crests) of a wave ○ Longer wavelength = lower frequency = lower energy ○ Shorter wavelength = higher frequency = higher energy Relationship of Energy, Wavelength, and Frequency  Frequency is directly proportional to energy ○ Gamma rays have highest frequency ○ For visible light, the color purple has the highest energy and red has the lowest energy Wavelength and frequency are indirectly proportional to each other Electrons, Quanta, and Flame Test Ground state The most stable state of an electron. Excited state The temporary high-energy position of an electron where it jumps to a higher energy level ○ When an electron absorbs light, it becomes excited. ○ When an electron emits light, it goes back to the ground state. The light emitted is in the form of colored light The same amount of energy is received and emitted by the atom Elements have different colors because of this  Niels Bohr's Explanation of Why Atoms Emit Light: The illustration on the left shows an atom absorbing light (the squiggly thing) and getting excited, and the one on the right shows the electron emitting the light and going back to the ground state Flame Test Used to identify metals and compounds Only metals can be identified as they lose electrons, so non-metals cannot be identified using this test Example: NaCl & NaI ○ Since Na is a metal, it will produce the light. Chlorine and Iodine are non-metals so they do not produce light. Compounds with the same metal element can produce similar light. Example: NaCl and KCl ○ Na and K are different metal elements, thus they will produce different light. 4 Quantum Numbers 1. Principal Quantum Number (n) Refers to the energy level It’s possible values are any numbers ≥1 (any positive value) 2. Secondary Quantum Number (l) Also known as azimuthal quantum numbers Refers to the sublevels: s, p, d, f  Sublevel Secondary Quantum No. of Electrons No. of Orbital Number s 0 2l 1 p 1 6l 3 d 2 10l 5 f 3 14l 7 3. Magnetic Spin Quantum Number (ml) The total number of orbitals in a sublevel For a given value l, the value of ml ranges between -l to +l Orbital is represented by ▢ ○ There are a maximum of 2 electrons per orbital Electrons are represented by ↑↓ 4. Electron Spin Quantum Number (ms) Gives insight into the direction in which the electron is spinning (clockwise & counter-clockwise) The only possible values: + ½, - ½ ○ The positive value of ms: upward spin ○ The negative value of ms: downward spin EXAMPLES Superscript Principal Secondary Magnetic Spin Electron Spin Quantum Quantum Quantum Quantum Number (n) Number (l) Number (ml) Number (ms) 1. 3s1 3 0 (since s = 0) 0 (since -0 to +0) +½ 2. 3p3 3 1 +1 (the orbitals +½ (upward stopped getting spin) filled at +1) 3. 3p4 3 1 -1 -½ 4. 4d5 4 2 +2 +½ Illustration of the Examples: 4.7 Periodic Trends (10/30/24) Recap: Periodic Law the physical & chemical properties of an element tend to vary periodically in order of increasing atomic number Periodic Trends 1. Atomic radius/size 2. Ionic radius/size 3. Ionization energy 4. Electron affinity 5. Electronegativity 6. Metallic property 7. Non-metallic property Atomic radius/size Distance from the nucleus of a neutral atom to the outermost shell Increases as you go from TOP to BOTTOM, ○ Higher periods mean higher energy levels that go farther from the nucleus Decreases from LEFT to RIGHT, ○ More electrons mean greater nuclear attraction, the electrons get closer to the nucleus Therefore, francium is the biggest and fluorine is the smallest  Note: francium is so big and useless (imagine only having 22 minutes to live bruh moment) Ionic radius/size Distance from the nucleus of an ion to the outermost electron orbital ○ Recap: Ions are charged atoms Cation is positively charged Anion is negatively charged Increases from TOP to BOTTOM, ○ Higher energy levels go farther from the nucleus Decreases from LEFT to RIGHT, ○ When atoms lose electrons, and the number of protons is still the same, the repulsion between electrons decreases and the attraction of the nucleus is stronger, therefore electrons will get pulled towards the nucleus and the size becomes smaller To put it in simple words, the electron cloud gets smaller when an atom loses electrons, so its size gets smaller ○ When atoms gain electrons, the number of protons remains the same, however, the repulsive force between the electrons becomes stronger than the attraction between the electrons and the nucleus, therefore expanding the electron cloud and increasing the size ○ This only applies to ions with the same type of charge (either cation or anion) because anions can be bigger than cations For example, Na+ has a bigger size than Mg2+, but they are both smaller than Cl- So Fe2+ has less ionic size than Fe Ionization energy The energy needed to remove an electron from an atom The ionization energy of the first electron removed is called the 1st ionization energy ○ The ionization energy of the second electron removed is called the 2nd ionization energy Increases from LEFT to RIGHT ○ As we go from left to right, the number of electrons in the valence shell increases and nuclear attraction becomes greater, so it becomes harder to remove an electron Why? For example, if you remove an electron from Na (1 valence electron in its neutral form) it becomes isoelectronic with a noble gas that has a stable configuration (8 valence electrons)  - We know that it would be hard to remove an electron from an atom if its in a stable configuration, because why would u even need to remove one. youre already stable!! ○ It decreases as you go from right to left, it requires lower energy to remove an electron because less electrons mean lower attraction between nucleus, so it becomes easier to remove electrons Decreases from TOP to BOTTOM, ○ Higher energy levels mean electrons farther from the nucleus and less nuclear attraction Ionization energy is opposite the atomic radius ○ Therefore, higher ionization energy means a smaller atomic radius ○ And lower ionization energy means a bigger atomic radius Electron Affinity Energy released when an atom accepts an electron to form an anion ○ Can be measured only in its gaseous state Increases from LEFT to RIGHT ○ It’s easier for an atom with more valence electrons to accept an electron because their electrons are closer to the nucleus. This means that the nucleus is strong enough to pull more electrons - It is also because they are closer to a stable configuration (noble gases) Decreases from TOP to BOTTOM ○ If you go from top to bottom, there will be more energy levels. The outermost energy level would be farther away from the nucleus, therefore it would be easier to lose an electron from the outermost energy level than to gain another as there is a lower attraction to the nucleus Electronegativity The ability of an atom to attract electrons and gain a negative charge  Increases from LEFT to RIGHT ○ If you go from left to right, it becomes easier for an atom to attract another electron as their configuration is near noble gas configuration ○ If you go from right to left, atoms tend to lose electrons instead of gaining ○ For example: Magnesium has 2 valence electrons. It would be easier for magnesium to lose its 2 valence electrons to attain noble gas configuration than to gain more electrons. Decreases from TOP to BOTTOM ○ Same reason as electron affinity Fluorine is the most electronegative element Metallic property The properties of a metal include: ○ Big atomic radius ○ Low ionization energy ○ Low electron affinity ○ Low electronegativity ○ Good conductor Increases as you go from TOP to BOTTOM Decreases as you go from LEFT to RIGHT Nonmetallic property The properties of a non-metal include: ○ Small atomic radius ○ High ionization energy ○ High electron affinity ○ High electronegativity ○ Good insulator Increases as you go from LEFT to RIGHT Decreases as you go from TOP to BOTTOM Top-Bottom: Increase, Left-Right: Decrease Top-Bottom: Decrease, Left-Right: Increase Atomic Size Ionization energy Ionic Size Electron affinity Metallic Property Electronegativity Nonmetallic property UNIT V: Stoichiometry Chemical Formulas and Ions Monatomic Ions to be Familiar With Family 1A 2A 3A 5A 6A 7A Charges +1 +2 +3 -3 -2 -1 Ions Li+ Be2+ B3+ N3- O2- F- (Fluoride) (Lithium (Beryllium (Boron ion) (Nitride) (Oxide) ion) ion) Na+ Mg2+ Al3+ P3- S2- Cl- (Sodium (Magnesium (Aluminum (Phosphide) (Sulfide) (Chloride) ion) ion) ion) K+ Ca2+ Br- (Potassium (Calcium (Bromide) ion) ion) I- (Iodide) Polyatomic Ions to be Familiar With Ion Formula Nitrite NO2- Nitrate NO3- Sulfite SO32- Sulfate SO42- Ammonium NH4+ Phosphate PO43- Phosphite PO33- Hydroxide OH- Hypochlorite ClO- Chlorite ClO2-  Chlorate ClO3- Perchlorate ClO4- Chromate CrO42- Dichromate Cr2O72- Permanganate MnO4- Acetate C2H3O2- Carbonate CO32- Bicarbonate HCO3- Cyanide CN- 5.1 Naming and Writing Ionic Formulas (11/12/24) Ionic compound compound composed of ions (cation + anion) Binary ionic compound: ○ ionic compound composed of two elements ○ Examples: Sodium chloride (NaCl), Potassium bromide (KBr), Zinc sulfide (ZnS) Steps in naming ionic compounds 1. Identify the cation and anion 2. Name the cation by element name (for example, a sodium cation would be named.. sodium) 3. Name anion by changing the last letters to -ide (for example, a fluorine anion would be fluoride) 4. Always write the cation first before the anion Example: Formula Cation Anion Name NaCl Na+ Cl- Sodium chloride MgO Mg2+ O2- Magnesium oxide Criss cross method Only applicable to ionic compounds The charge of one element will be the subscript/number of atoms of the other Examples: 1. Identify the charges of each atom 2. The charge of sodium will determine the number of chlorine atoms, and the charge of chlorine will determine the number of sodium atoms a. So there’s only 1 sodium and 1 chlorine atom as sodium has +1 and chlorine has -1 1. Identify the charges of each atom 2. The charge of magnesium will determine the number of fluorine atoms, and the charge of fluorine will determine the number of magnesium atoms a. Since magnesium has a charge of 2+ there would be 2 fluorine atoms, and since fluorine has a charge of -1, there would be 1 magnesium atom  theyre meant to be Transition Metals form Cations Most transition metals and Group 4A metals form positive ions  Unlike representative elements, whose charges are fixed, the charges of transition metals vary Charges of Transition and Group 4A Metals (not all of these are varying) Transition Metals Group 4A Cr2+ Sn2+ Cr3+ Sn4+ Fe2+ Pb2+ Fe3+ Pb4+ Cu+ Cu2+ Ag+ Au+ Au3+ Zn2+ Cd2+ Stock system Metals that form more than 1 cation use a Roman numeral to identify ionic charge Used for elements that have varying charges ○ guys pls dont do the same mistake i made i used stock system on yttrium and zirconium and silver grgrgrgr Example: ○ Cr2+ - Chromium (II) ○ Cr3+ - Chromium (III) Classical system Uses the Latin names of elements of varying charges ○ Use the suffix “-ic” for higher charge Example: Fe3+ - Ferric ○ Use the suffix “-ous” for lower charge Example: Fe2+ - Ferrous 5.2 Naming and Writing Covalent and Molecular Formulas (11/20/24) Covalent Compounds  Also called molecular compounds Compounds composed of two or more non-metals ○ Example: Water (H2O), Carbon dioxide (CO2) Steps in Naming Covalent Compounds 1. Like cations, first element keeps its name, but you will have to add prefixes if there are 2 or more atoms of the element a. Example: H2 - Dihydrogen b. Instead of saying Monocarbon Monoxide, we just say Carbon Monoxide (CO) 2. Like anions, we change the last letters of the second element into -ide, and we have to add prefixes Prefixes in Writing Covalent Compounds 1. One = mono- (only for second element) 2. Two = di- 3. Three = tri- 4. Four = tetra- 5. Five = penta- 6. Six = hexa- 7. Seven = hepta- 8. Eight = octa- 9. Nine = nona- 10. Ten = deca- Examples H2O - Dihydrogen Monoxide SF6 - Sulfur Hexafluoride NO2 - Nitrogen Dioxide Note: There’s a rule when it comes to putting prefixes on oxide. Instead of saying “tetraoxide”, we say “tetroxide”. (We drop the -a of the prefix) Acids Compounds made of a hydrogen ion and a nonmetal or anion ○ Binary acids - contain hydrogen and a nonmetal ○ Ternary acids - also called oxyacids, acids that contains oxygen Naming Binary Acids Formula Name as pure compound Name in aqueous solution HCl Hydrogen chloride Hydrochloric acid HF Hydrogen fluoride Hydrofluoric acid Naming Ternary Acids Formula Name as pure compound Name in aqueous solution  H2SO4 Hydrogen sulfate Sulfuric acid H2SO3 Hydrogen sulfite Sulfurous acid Bases Compounds made of a hydroxide anion and a metal Note: Most bases are ionic compounds, as they are composed of metals and nonmetals, so they are not covalent (idk why maam included this bc the discussion is about covalent compounds but alr) Naming Bases (They are usually named the same way as ionic compounds) Formula Name NaOH Sodium hydroxide Ca(OH)2 Calcium hydroxide 5.3 Balancing Chemical Equations (11/26/24) Chemical equation The symbolic representation of a chemical reaction in the form of symbols and formulas Two parts: ○ Reactants (left side) ○ Products (right side) List of Diatomic Molecules (gases) Symbol Name  H2 Hydrogen gas F2 Fluorine gas Cl2 Chlorine gas Br2 Bromine gas I2 Iodine gas O2 Oxygen gas N2 Nitrogen gas Symbols used in Chemical Equations (s) Solid (l) Liquid (g) Gas (aq) Aqueous, dissolved in water ↑ Gas forms ↓ Solid precipitate (insoluble solid) forms -> ”To yield/produce” Reaction in which products can reform into ⇌ reactants (Reversible reaction) Heat is applied Other conditions (temperature, pressure, etc.) Why do we balance? Lavoisier’s Law of Conservation of Mass Before, during, and after a chemical reaction in a closed system, matter is neither created nor destroyed and mass remains constant. ○ To put in simpler words, mass of reactants = mass of products Key Notes Subscripts cannot be changed, only the coefficients. ○ If we change the subscript of a chemical, it would be considered incorrect as that would produce another different chemical. We read the coefficients as moles. ○ For example, 2Na is read as “2 moles of sodium” The reaction “A + B -> C + D” is read as “1 mole of A reacts with 1 mole of B to yield/produce 1 mole of C and 1 mole of D”. Steps in Balancing Chemical Equations 1. Count the number of atoms on each side. 2. Ensure the number of atoms for each element is equal on both the reactants and products sides. 3. Add coefficients to make the number of atoms equal using trial and error. ○ There’s no other way to balance equations other than trial and error, but practicing will help u improve Q3 9-CONCHEM Q3 - Chemistry Ma’am Mary Edd Anne Certeza | 3rd Quarter | Owned by Ysabelle Ng, Lindsay Inocando, Stephen Villarin, Stephanie Trinidad Q4 9-CONCHEM Q4 - Chemistry Ma’am Mary Edd Anne Certeza | 4th Quarter | Owned by Ysabelle Ng, Lindsay Inocando, Stephen Villarin, Stephanie Trinidad