General Chemistry 1 Midterm Reviewer PDF

GENERAL CHEMISTRY 1 CHEM1-MID FIRST SEMESTER | ACADEMIC YEAR 2024-2025 | MR. MARK BENEDICK TOPINIO IX.​ The Mole Concept...

GENERAL CHEMISTRY 1 CHEM1-MID FIRST SEMESTER | ACADEMIC YEAR 2024-2025 | MR. MARK BENEDICK TOPINIO IX.​ The Mole Concept COURSE OUTLINE a.​ Molar Mass X.​ Percent Composition I.​ Introduction To Chemistry a.​ Steps to Solve a.​ What is chemistry b.​ Tips to Solve b.​ Major Branches of Chemistry c.​ Example With The Use Of Steps c.​ Minor Branches of Chemistry XI.​ Empirical Formula and Molecular Formula d.​ The Scientific Method a.​ Steps to Solve (Empirical Formula) e.​ 3 Types of Variables b.​ Steps to Solve (Molecular Formula) II.​ Matter and Its Properties XII.​ Sample Problem Questions a.​ Properties and Classification of Matter XIII.​ Answer Keys b.​ Physical Properties and Chemical Properties of Matter c.​ Methods in Separating Components of INTRODUCTION TO CHEMISTRY Mixtures What is Chemistry III.​ Measurements ➔​ A branch of science concerned with the properties, a.​ Scientific Measurements b.​ Conversion of Units and Dimensional Analysis composition, and structure of substances and the c.​ Significant Figures and Scientific Notations changes they undergo. d.​ Accuracy and Precision ➔​ It has been called as the Central Science. e.​ Density and Specific Gravity ➔​ It is the study of matter and changes it undergoes. IV.​ Atomic Theory & Basic Laws of Matter a.​ Atomic Theory b.​ Demodcritus (400 BC) Major Branches Of Chemistry c.​ Aristotle 1.​ Analytical Chemistry d.​ John Dalton (1766-1844) ➔​ Material samples are analyzed to gain an e.​ Joseph John Dalton (1859-1940) understanding of their chemical composition. f.​ Ernest Rutherford 2.​ Biochemistry g.​ Eugene Goldstein ➔​ Living organisms are studied to understand the h.​ James Chadwick (1891-1974) chemical composition, reactions, and i.​ Niels Bohr j.​ Erwin Schrodinger interactions that occur within the organism. k.​ Model Limitations 3.​ Inorganic Chemistry l.​ Fundamental Laws of Matter ➔​ Inorganic compounds are studied to understand m.​ Antoine – Laurent de Lavoisier (1743-1794) their chemical compositions, properties, and n.​ Joseph-Louis Proust (1753-1826) reactions that occur within the sample. V.​ Atomic Structure & Subatomic Particles 4.​ Organic Chemistry a.​ Atoms b.​ Subatomic Particles ➔​ Organic compounds are studied to understand c.​ Chemical Notation of Atoms their chemical compositions, properties, and d.​ Isotopes reactions that occur within the sample. e.​ Atomic Average Mass 5.​ Physical Chemistry VI.​ Molecules and Ions ➔​ The physical processes and general a.​ Molecules characteristics of chemical systems are studied b.​ Ions to understand the energy levels and chemical VII.​ Naming Chemical Compounds a.​ Chemical Formula transformations that occur. b.​ Binary Ionic Compounds c.​ Ionic Compounds with Polyatomic Anions Minor Branches Of Chemistry d.​ Ionic Compounds with Variable Valence ➔​ Nuclear, Environmental, Material, Forensic, Geo, e.​ Binary Covalent Compounds Medicinal, Clinical, Organometallic, Polymer, Stereo, Solid f.​ Naming Acids g.​ Supplemental Videos and Websites state, Bioinorganic, Bioanalytical, Photochem, Surface, VIII.​ Balancing Chemical Equation Quantum, Spectroscopy, Molecular Biology, Genetics, a.​ Stoichiometry Enzymology, Endocrinology, Astro, Food, Cluster, b.​ Chemical Equations Combinatorial, Green, Theoretical and many more. c.​ Steps in Balancing Chemical Equations The Scientific Method General Chemistry 1: MIDTERM NOTRE DAME OF GREATER MANILA — ACADEMIC SCHOOL YEAR 2024-2025 1  GENERAL CHEMISTRY 1 CHEM1-MID FIRST SEMESTER | ACADEMIC YEAR 2024-2025 | MR. MARK BENEDICK TOPINIO 1.​ Make and observation and identify the problem ➔​ No definite volume and shape. ➔​ Observe a phenomenon and define the specific ➔​ Widely separated by great distances. problem or question. ➔​ Just like liquids, gases can flow. 2.​ Formulate a hypothesis 4.​ Plasma ➔​ Propose a testable explanation or prediction ➔​ Composed of charged particles. (+, -). based on prior knowledge. ➔​ Exhibit properties distinct from ordinary gases, 3.​ Gather the Data information liquid or solids. ➔​ Collect background information and resources related to the problem. Changes of Matter 4.​ Conduct and experiment 1.​ Melting - a physical process that results in the phase ➔​ Design and perform an experiment to test the transition of a substance from a solid to a liquid. hypothesis under controlled conditions. 2.​ Freezing - a phase transition in which a liquid turns into a 5.​ Collect and analyze the data results solid when its temperature is lowered below its freezing ➔​ Record observations and data, then analyze point. them to see if they support the hypothesis. 3.​ Vaporization - conversion of a substance from the liquid or 6.​ Draw and make a conclusion solid phase into the gaseous (vapour) phase. ➔​ Summarize the findings and determine whether 4.​ Condensation - is the change of the state of matter from the hypothesis was supported or refuted. the gas phase into the liquid phase, and is the reverse of 7.​ Publish and/or communicate your research study vaporization ➔​ SHre the results with scientific community 5.​ Sublimation - conversion of a substance from the solid to through publications or presentations. the gaseous state without its becoming liquid. 6.​ Deposition - is the transition of a substance directly from 3 Types Of Variables the gas to the solid state on cooling, without passing 1.​ Dependent Variable through the liquid state. ➔​ The change thet happens because of the 7.​ Evaporation - is the process that changes liquid water to independent variable. gaseous water (water vapor) 2.​ Independent Variable 8.​ Recombination - the transition from plasma to gaseous ➔​ The one thing you change. Limit to only one in state occurs if electrons and ions are left to recombine. an experiment. 9.​ Ionization - the transition from the gaseous to the plasma 3.​ Controlled Variable state. ➔​ Everything you want to remain constant and unchanging. Classification of Matter MATTER AND ITS PROPERTIES Properties and Classification of Matter What is Matter ➔​ It is anything that occupies space and has mass. ➔​ All matter is composed of very tiny particles called atoms. States of Matter 1.​ Solids ➔​ Definite shape and volume ➔​ Tightly packed ➔​ It has a little freedom of motion 2.​ Liquids 1.​ Pure Substance - type of matter that is composed of only ➔​ Definite volume but no definite shape. one kind of substance, one kind of atom or one kind of ➔​ Loosely packed. molecule. They are homogeneous or one-phase in ➔​ It has a little spaces enough for the appearance. molecules/atoms to pass one another. a.​ Elements - type of pure substances that are 3.​ Gas made up of one kind of atom. Examples are gold, iron, carbon, silicon and helium. General Chemistry 1: MIDTERM NOTRE DAME OF GREATER MANILA — ACADEMIC SCHOOL YEAR 2024-2025 2  GENERAL CHEMISTRY 1 CHEM1-MID FIRST SEMESTER | ACADEMIC YEAR 2024-2025 | MR. MARK BENEDICK TOPINIO b.​ Heterogenous Mixture - type of mixtures that consists of two or more distinct proportions and Metals Non-Metals Metalloids or compositions, which can be separated Semi-Metals mechanically from one another. Elements Elements Elements Colloid Suspension characterize which are which d by their generally possesses brilliant light without the It possess Tyndall A heterogenous luster, brilliant characteristi effect with Brownian mixture of two or ductility, luster, brittle cs of metals motion movement more substances in malleability and are poor and non- and particles are which the particles and good conductor of metals. suspended in a are much larger than conductors heat and medium. colloids and will of heat and electricity. settle out on electricity. standing. b.​ Compounds - type of pure substances that are Comparison on the Different Kinds of Mixture chemical combinations of elements in a fixed ratio. Examples are table salt (sodium chloride, Solutions Colloids Suspensions NaCl), table sugar and water (H2O). Homogenous Heterogenous Heterogenous Organic Compound Inorganic Compound Particle size: Partilce size: Particle size: over 0.01-0.1nm; atoms, 1-1000 nm, 1000 nm, Produced by living Produced by ions, or molecules dispersed; large suspended; large things. Atoms are non-living natural molecules or particles or held by covalent processes. Atoms are aggregates aggregates bond. CAPLIN held by Ionic bond. (Carbohydrates, BOSWAG (Bases, Do not separate on Do not separate on Particles settle out Proteins, Lipids, and Oxides, Salts, Water, standing standing Nucleic Acids) Acids, and Gases) Cannot be Cannot be Can be separated separated by separated by by filtration 2.​ Mixture - type of matter that is composed of several filtration filtration substances, either physical combinations of compounds or Do not scatter light Scatter light May either scatter elements or It is produced by mixing two or more pure (Tyndall Effect) light or be opaque substances. a.​ Homogenous Mixture - type of mixtures that have one-phase appearance. You cannot easily Physical Properties and Chemical Properties distinguish the components. Properties of Matter ➔​ It generally pertain to the characteristics of a certain substance that sets it apart from other substances. Solutions A system or a homogenous mixture in which one or more ➔​ It can be categorized into two groups – the Physical substances are homogenously mixed or dissolved to other properties and the Chemical Properties. substance. Solvent Solute Physical Properties ➔​ The physical properties of matter are those that can be The dissolving agent The component that measured and observed without changing the or the most is dissolved or the composition of the substance. abundant least abundant component. component in the solution. Intensive Property Extensive Property General Chemistry 1: MIDTERM NOTRE DAME OF GREATER MANILA — ACADEMIC SCHOOL YEAR 2024-2025 3  GENERAL CHEMISTRY 1 CHEM1-MID FIRST SEMESTER | ACADEMIC YEAR 2024-2025 | MR. MARK BENEDICK TOPINIO ➔​ Measurement is a process of expressing size or quantity Do not depend on how much Depend on how much matter matter is being considered. is being considered. of same kind into a known standard unit. ➔​ Observations Examples are Density, Examples are Mass, Length, a.​ Quantitative Observations Temperature, Viscosity, Volume, Width, Area, Height, ​ Measurement Melting point, Boiling point, and Surface area. ​ Number and a unit (scale) Ductility and Malleability. ​ Both parts must be present in order for the measurement to be Chemical Properties meaningful. ➔​ The chemical property of matter is the ability of a b.​ Qualitative Observations substance to react with other substances such as air, ​ Does not involve a number. water, acid, and base. ➔​ Two things to be considered about the product of ➔​ Combustibility, Stability, Reactivity, Relative, Activity, measurement: Ionization, Toxicity 1.​ Quantity expressed in number 2.​ The proper unit ➔​ Standard Systems of Units: 1.​ English System 2.​ Metric System The System of Measurements 1.​ SI Units - Le système International d'unités (French Phrase) ​ Meter (m) – as a unit of length ​ Kilogram (kg) – as a unit of mass ​ Second (s) – as a unit of time Physical and Chemical Changes 2.​ CGS System - centimetre–gram–second system of units. Matter undergoes changes when its conditions are altered. These ​ Centimeter (cm) – as a unit of length changes can either be physical or chemical changes. ​ Gram (g) – as a unit of mass ​ Second (s) – as a unit of time 1.​ Physical Changes - matter changes form but not chemical 3.​ British Engineering (BE) System – The gravitational version identity. system. a.​ Examples: Breaking glass, Shredding paper, ​ Foot (ft) – as a unit of length Folding paper, Melting ice, Boiling water, Cutting ​ Slug (sl) – as a unit of mass hair, Cutting wood, Ice sublimation, Mixing ​ Pound (lb) – as a unit of weight candies, Breaking eggs. ​ Second (s) – as a unit of time 2.​ Chemical Changes - a chemical reaction occurs and new Two Kinds of Units / Quantities products are formed. 1.​ Base (Fundamental) Units / Quantities a.​ Examples: Burning wood, Digestion, Rotting ➔​ The word base refers to the fact that these units fruit, Souring milk, Baking cake, Exploading (length, mass, time, electric current, fireworks, Rusting metals, Cooking an egg, Using temperature, luminous intensity, and amount of a battery, Photosynthesis. substance) are used along with various laws to b.​ Evidences of chemical change: define additional units for other important 1.​ A change in color, odor and taste. physical quantities such as force and energy. 2.​ Formation of precipitate in a chemical reaction. 3.​ Release of Heat or Absorption of heat. 4.​ New substances are formed. MEASUREMENTS Scientific Measurements General Chemistry 1: MIDTERM NOTRE DAME OF GREATER MANILA — ACADEMIC SCHOOL YEAR 2024-2025 4  GENERAL CHEMISTRY 1 CHEM1-MID FIRST SEMESTER | ACADEMIC YEAR 2024-2025 | MR. MARK BENEDICK TOPINIO Given Unit Dimensional Unit 2.​ Derived Units / Quantities ➔​ They are combinations of the base Given Unit units/quantities. ➔​ Examples: 𝑑 𝐷𝑖𝑓𝑓𝑒𝑟𝑒𝑛𝑡 𝑈𝑛𝑖𝑡 Velocity: 𝑣 = 𝑡 𝑆𝑎𝑚𝑒 𝑈𝑛𝑖𝑡 𝑥 𝑆𝑎𝑚𝑒 𝑈𝑛𝑖𝑡 Force: 𝐹 = 𝑚𝑎 Potential Energy: 𝑃𝐸 = 𝑚𝑔ℎ 𝐼𝑛𝑡𝑒𝑟𝑚𝑒𝑑𝑖𝑎𝑡𝑒 𝑈𝑛𝑖𝑡 𝐷𝑒𝑠𝑖𝑟𝑒𝑑 𝑈𝑛𝑖𝑡 𝐺𝑖𝑣𝑒𝑛 𝑈𝑛𝑖𝑡 𝑥 𝐺𝑖𝑣𝑒𝑛 𝑈𝑛𝑖𝑡 𝑥 𝐼𝑛𝑡𝑒𝑟𝑚𝑒𝑑𝑖𝑎𝑡𝑒 𝑈𝑛𝑖𝑡 Mass Weight Density Volume amount or force that compactness amount of ➔​ When doing dimensional analysis problems, follow this quantity of gravity exerts of matter; space list of steps: matter on an object mass per unit occupied by 1.​ Identify the given. volume matter 2.​ Identify conversion factors that will help you get kilogram (kg) W=mg Newton (N) (ρ =m/v) Cubic meter from your original units to your desired unit. or kg (m/s2) kg/m3 (cm3) 3.​ Set up your equation so that your given units cancel out to give the desired units. A unit will cancel out if it appears in both the numerator Conversion of Units and Dimensional Analysis and the denominator during the equation. Conversion of Units 4.​ Multiply through to get your final answer. Don’t ➔​ Conversion is a process of expressing one quantity from forget the units and the significant figures. its original unit to a desired unit. ➔​ Sometimes, we must convert units from one system to Significant Figures and Scientific Notations another. This will allow us to use different quantities in Significant Figures problems involving a different unit. In order to do this, a ➔​ These are the number of digits in a value, usually a conversion factor is used. measurement, that contribute to the value’s accuracy. ➔​ A conversion factor serves as a ratio of two different units ➔​ The number of digits in a given value or a measurement, to attain a desired unit in a systematic approach. necessary to decide the accuracy and precision of measurement. Going to the left, move the decimal point ot the left. Going to the right, move the decimal point to the right. Dimensional Analysis Rules of Significant Figures ➔​ Also called Factor Label Method or Unit Factor Method 1.​ All numbers that are nonzero digits in a measurement are because we use conversion factors to get the same units. significant. (3999.79 has six significant figures) General Chemistry 1: MIDTERM NOTRE DAME OF GREATER MANILA — ACADEMIC SCHOOL YEAR 2024-2025 5  GENERAL CHEMISTRY 1 CHEM1-MID FIRST SEMESTER | ACADEMIC YEAR 2024-2025 | MR. MARK BENEDICK TOPINIO 2.​ Zeros that appear between other nonzero digits or captive 1.​ Whenever our bases are the same and are being zeros are always significant. (1.80205 has six significant multiplied, we add the exponents together. ((2.6 x 105 ) figures) (3.7 x 104 )). 3.​ Zeros after the decimal are significant. (43.000 has five 2.​ Whenever our bases are the same and are being divided, 2 significant figure) we subtract the exponents together. 10 −4 4.​ Zeros and coefficient in scientific notations are significant. 10 (6.022 x 1023 has four significant figure) 3.​ Whenever our base is raised to another power or 5 3 5.​ Leading or beginning zeroes are NOT significant. (0.0004 exponent, we will multiply the exponents together. (10 ) has one significant figure) 6.​ Zeros in a large number without a decimal are not Accuracy and Precision significant. (7,000 has one significant figure) 7.​ When ADDING or SUBTRACTING a number, look for the LEAST number of decimal places. A complete measurement has ➔​ 22. 13 + 17. 0 + 2. 024 = 41. 154 ≈ 41, 2 Number or Unit Degree of ➔​ 5365. 999 – 234. 66706 = 5131. 33194 = 5131. 332 Magnitude Reliablity 8.​ When MULTIPLYING or DIVIDING a number, look for the LEAST number of significant figures. ➔​ 24. 5 𝑥 63. 2751 = 1, 550. 23995 ≈ 1, 550 ➔​ Accuracy refers to how close a measured value is to the 1000.1 real or “true” value ➔​ 243 = 4. 11563786 ≈ 4. 12 ➔​ Precision refers to the degree of reproducibility of a measured quantity or how close the measurements are to Scientific Notations each other. ➔​ Scientific notation is a form of presenting very large numbers or very small numbers in a simpler form. Types of Errors 1.​ Random Error - means that a measurement has an equal Standard Notation Scientific Notation probability of being high or low. ​ Example: Blood Pressure 0.00000524 −6 5, 24 𝑥 10 2.​ Systematic Error - it occurs in the same direction each time; it either always high or always low. 320,000,000 8 3. 2 𝑥 10 ​ Example: Temperature measured using thermal ➔​ How to determine if the exponent will be negative or scanner. positive? 1.​ Negative Exponent Density and Specific Gravity ➔​ All small numbers, which are numbers 1.​ Density less than 1, have negative exponents. ➔​ It is usually expressed as grams per cubic ➔​ 𝑁𝑢𝑚𝑏𝑒𝑟𝑠 < 1 centimeter. 2.​ Positive Exponent ➔​ The density of the substance is the relationship ➔​ All large numbers, which are numbers between the mass of the substance and how greater than or equal to 1, have much space it takes up volume. 𝑚𝑎𝑠𝑠 (𝑚) positive exponents. ➔​ Formula: ρ = 𝑣𝑜𝑙𝑢𝑚𝑒 (𝑣) ➔​ 𝑁𝑢𝑚𝑏𝑒𝑟𝑠 ≥ 1 ➔​ One can say that an object is denser when it 3.​ Exponents of Zero sinks, while it is less dense when it floats. ➔​ Numbers that already have one ➔​ For regularly shaped solid, we can determine number that is not zero followed by a the volume by measuring its applicable decimal. dimensions and using the appropriate formula. ➔​ For irregularly shaped solid, determine its volume using the water displacement method. Performing Operations (Rules of Exponents) General Chemistry 1: MIDTERM NOTRE DAME OF GREATER MANILA — ACADEMIC SCHOOL YEAR 2024-2025 6  GENERAL CHEMISTRY 1 CHEM1-MID FIRST SEMESTER | ACADEMIC YEAR 2024-2025 | MR. MARK BENEDICK TOPINIO Atomic Theory ➔​ Atomic theory is a scientific concept that explains the nature of matter and its behavior. ➔​ Several theories try to explore the structure of atoms and their interactions with other atoms and molecules. ➔​ Atomic theory has evolved over centuries with the contributions of many scientists. Democritus (400 BC) ​ A Greek philosopher, the first person to think about an atom’s existence. ​ His theory: Matter could not be divided into smaller and smaller pieces forever, eventually the smallest possible piece would be obtained. 𝑤ℎ𝑒𝑟𝑒: ​ He believed that matter was composed of tiny indivisible 𝑀 = 𝑚𝑎𝑠𝑠 𝐷 = 𝑑𝑒𝑛𝑠𝑖𝑡𝑦 particles, he called it atomos, meaning “not to be cut”. 𝑉 = 𝑣𝑜𝑙𝑢𝑚𝑒 ​ He had no experimental evidence to support this thoughts. 2.​ Specific Gravity ➔​ The ratio of the density (mass of a unit volume) Aristotle of a substance to the density (mass of the same ​ He did not believe in the existence of indivisible atoms. unit volume) of a reference substance chosen as He rejected that matter was composed of tiny, indivisible a standard, both having the same temperature. particles. ➔​ Water is used as the standard for the specific ​ He proposed that all matter was made up of four gravities of liquids and solids. fundamental elements: earth, water, air, and fire. ➔​ S.G. is unitless. ​ The relationships among these elements and the four principles – hot, moist, dry, and cold. 𝐷𝑒𝑛𝑠𝑖𝑡𝑦 𝑜𝑓 𝑂𝑏𝑗𝑒𝑐𝑡 ​ The fifth element proposed by Aristotle is the Aether. 𝑆𝑝𝑒𝑐𝑖𝑓𝑖𝑐 𝐺𝑟𝑎𝑣𝑖𝑡𝑦 = 𝐷𝑒𝑛𝑠𝑖𝑡𝑦 𝑜𝑓 𝑤𝑎𝑡𝑒𝑟 John Dalton (1766-1844) ➔​ When we compute specific gravity, we're ​ The English chemist, performed a number of experiments determining how the density of a substance that eventually led to the acceptance of the idea of compares to the density of a reference atoms. substance, typically water (for liquids and solids) ​ He described atoms as tiny invisible particles that could or air (for gases). not be created, destroyed, or divided. ➔​ Specific Gravity = 1.0: If a substance has a ​ Dalton had four major points (postulates) to his theory. specific gravity of 1.0, it means its density is exactly the same as the density of the reference Dalton’s Theory substance (water, if we're using that as the 1.​ All elements are composed of indivisible particles called reference). atoms. ➔​ Specific Gravity > 1.0: If a substance has a 2.​ Atoms of the same element are identical. The atoms of specific gravity greater than 1.0, like 2.3, it any one element are different from those of another. means the substance is denser than the 3.​ Atoms of different elements mix or combine in whole reference. A specific gravity of 2.3 means the number ratios. Example: Oxygen combines with hydrogen substance is 2.3 times denser than water. to form water in a 2:1 ratio. ➔​ Specific Gravity < 1.0: If a substance has a 4.​ Chemical reactions occur when atoms separate, join, or specific gravity less than 1.0, it means the rearrange. In a chemical reaction, atoms of one element substance is less dense than the reference NEVER change into another. substance. Billiard Ball Model Atomic Theory & Basic Laws of Matter General Chemistry 1: MIDTERM NOTRE DAME OF GREATER MANILA — ACADEMIC SCHOOL YEAR 2024-2025 7  GENERAL CHEMISTRY 1 CHEM1-MID FIRST SEMESTER | ACADEMIC YEAR 2024-2025 | MR. MARK BENEDICK TOPINIO ​ Proposed by John Dalton in 1804. cathode rays, these were later called canal rays or positive ​ It says that matter was composed of small, spherical rays. particles. ​ His work provided evidence for the existence of positively ​ He deduced that all elements are composed of atoms. charged particles in atoms, supporting the later Atoms are indivisible and indestructible particles. development of the nuclear model of the atom. Joseph John Thomson (1859-1940) James Chadwick (1891-1974) ​ In 1897, the English scientist, Thomson provided the first ​ He was a British physicist who made significant hint that an atom is made of even smaller particles. contributions particularly with his discovery of the ​ J.J. Thomson discovered the electron. neutron. ​ From his experimental evidence, he believed that the ​ Detected a neutral radiation that was not deflected by atom was a solid positive sphere with electrons shoved electric or magnetic fields. into the sides of it. ​ He proposed that neutrons are neutral particles located in the nucleus of an atom and have a mass about that of Plum Pudding Model a proton. ​ Atoms were made from a positively charged substance ​ His discovery explained isotopes and led to the with negatively charged electrons scattered about, like development of nuclear physics and atomic models. raisins in a pudding. ​ Also used cathode ray experiment to discover the Niels Bohr existence of the electron. ​ In 1913, the Danish scientist Niels Bohr proposed an improvement. In his model, he placed each electron in a Cathode Ray Tube Experiment specific energy level. ​ In the tube was an inert gas, and two plates, a positive and ​ Discovered that electrons exist in several distinct layers or a negative. levels. ​ The particles in the gas were attracted to the positive ​ Travel around nucleus like planets travel around sun plate. (Electrons Orbit). ​ Therefore, the particles MUST have a negative charge. ​ Electrons can jump between levels with energy being (Opposite Attracts) added/released. Ernest Rutherford Bohr Model ​ Discovered the nucleus and credited for the discovery of ​ Niels Bohr proposed that electrons revolve around the protons. central positive nucleus (like planets in the solar system). ​ Atoms are made of mostly empty space. ​ Bohr also suggested that the electrons can only revolve in ​ Used the gold foil experiment to discover the nucleus. certain orbits, or at certain energy levels (ie, the energy ​ Shot high energy beam of alpha particles into gold foil. levels are quantized). Nuclear Model The Wave Model ​ Rutherford found that most (99%) of the alpha particles ​ According to the theory of wave mechanics, electrons do that he shot at the gold went straight through, indicating not move about an atom in a definite path, like the that the atom is mostly empty space. planets around the sun. ​ From these experiments Rutherford concluded that the ​ In fact, it is impossible to determine the exact location of atom had a dense positive core, with the rest composed an electron. The probable location of an electron is based of mostly empty space with the occasional negatively on how much energy the electron has. charged electron. ​ According to the modern atomic model, at atom has a small positively charged nucleus surrounded by a large region in which there are enough electrons to make an Eugene Goldstein atom neutral. ​ German physicist known for the discovery of proton. ​ He conducted experiments using cathode ray tube, where he observed rays traveling in the opposite direction to Erwin Schrödinger General Chemistry 1: MIDTERM NOTRE DAME OF GREATER MANILA — ACADEMIC SCHOOL YEAR 2024-2025 8  GENERAL CHEMISTRY 1 CHEM1-MID FIRST SEMESTER | ACADEMIC YEAR 2024-2025 | MR. MARK BENEDICK TOPINIO ​ An Austrian physicist who developed the Schrödinger by Ernest nucleus against electrostatic repulsion equation and introduced the concept of wave mechanics Rutherford ​ Lacked details about electron orbits and in quantum theory. energy levels ​ Found that Electrons live in fuzzy regions or “clouds” not ​ Didn't incorporate the principles of quantum mechanics distinct orbits. ​ He proposed the Quantum Mechanical Model or Electron Quantum Model ​ Mathematical complexity of the model. Cloud Model. or Electron ​ Requires advanced mathematics to ​ Electrons are not confined to fixed energy levels; rather Cloud Model calculate electron probabilities. by Erwin ​ Doesn't provide a simple visual they occupy volumes of space outside the nucleus; representation of atomic structure. Schrodinger electron energy is based on its location (increases away from nucleus). Electron Cloud Model Fundamental Laws of Matter ​ A space in which electrons are likely to be found. 1.​ Law of Conservation of Mass (Lavoisier) ​ Electrons whirl about the nucleus billions of times in one ​ Mass is neither be created nor destroyed in a second. chemical reaction. ​ They are not moving around in random patterns. 2.​ Law of Definite Proportion (Proust) ​ Location of electrons depends upon how much energy the ​ A given compound always contains exactly electron has. same proportion of elements by mass. 3.​ Law of Multiple Proportions (Dalton) ​ When two elements form a series of Quantum Mechanical Model compounds, the ratios of the masses of the ​ This model sees the electrons not as individual particles, second element that combine with 1 gram of the but as behaving like a cloud - the electron can be first element can always be reduced to small “anywhere” in a certain energy level. whole numbers. ​ The electron cloud model is also known as the quantum mechanical model or quantum model of the atom. ​ This model represents the probability distributions of Antoine – Laurent de Lavoisier (1743 – 1794) electrons around the nucleus, rather than specific orbits, ​ French Chemist-Physicist and the Father of Modern and is based on the principles of quantum mechanics. Chemistry. ​ He had the first version of the Law of Conservation of Matter. Model Limitations ​ He also named oxygen and hydrogen. ​ He invented the first periodic table which had 33 Nuclear Model ​ Couldn’t explain differences in atomic elements. by Ernest mass within an element (isotopes) Rutherford ​ Didn’t account for the presence of subatomic particles like protons, Law of Conservation of Matter neutrons, and electrons. ​ The total weight of the substances entering into a ​ Couldn’t explain the behavior of atoms in chemical reactions. chemical change is equal to the total weight of the substances produced. Plum Pudding ​ Couldn’t explain why electrons didn’t ​ Mass is neither created nor destroyed, they’re just Model collapse into the positive sphere. rearranged. by J.J. Thomson ​ Failed to predict the distribution and arrangement of electrons. ​ Lacked explanation for the nucleus and Combustion Reaction its positive charge. Fuel + Oxygen Gas → Carbon dioxide + Water Vapor CxHy + O2 Planetary Model ​ Limited to explaining the hydrogen → CO2 + H2O by Neils Bohr atom. ​ Couldn't account for the behavior of multi-electron atoms. ​ In ordinary chemical reactions, the total mass of the ​ Didn't incorporate the wave-like nature reactants is EQUAL to the total mass of the products. of electrons. Nuclear Model ​ Didn't explain the stability of the General Chemistry 1: MIDTERM NOTRE DAME OF GREATER MANILA — ACADEMIC SCHOOL YEAR 2024-2025 9  GENERAL CHEMISTRY 1 CHEM1-MID FIRST SEMESTER | ACADEMIC YEAR 2024-2025 | MR. MARK BENEDICK TOPINIO ​ The total substances present at the end of chemical 2.​ Salt, or NaCl, is composed of Na and Cl atoms. For it to be process is the SAME as total mass of substances present created, both the sodium and the chlorine atoms have to be before process took place. in the same proportion. 3.​ Glucose is always made up of 6 carbon atoms, 12 hydrogen EXAMPLES atoms, and 6 oxygen atoms. This is the only ratio of these three elements that can make glucose. 1.​ If you form water with two (2) grams of hydrogen and sixteen 4.​ A 100g of the compound Carbon dioxide (CO2) always (16) grams of oxygen, the hydrogen and oxygen atoms are contains the two elements carbon and oxygen in the same still present in water and thus, a total weight of eighteen (18) proportion by mass which is 27.29 parts of carbon and 72.21 grams will still be present when water is formed. parts of oxygen gas. 2.​ The operation of the ordinary flashbulb. Fine wires of magnesium (metal) and oxygen gas are sealed within the bulb. When these reactants are energized, they combine chemically and produce magnesium oxide, a blinding white light and considerable heat. FORMULA 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑒𝑙𝑒𝑚𝑒𝑛𝑡 % 𝑜𝑓 𝑒𝑎𝑐ℎ 𝑒𝑙𝑒𝑚𝑒𝑛𝑡 = 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑐𝑜𝑚𝑝𝑜𝑢𝑛𝑡 3.​ The reaction of baking soda and vinegar. At first glance, the reaction seems to finish with less mass than it started with. But by placing a balloon atop the reaction container, one can 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑒𝑙𝑒𝑚𝑒𝑛𝑡∗𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑢𝑏𝑠𝑡𝑎𝑛𝑐𝑒 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑒𝑙𝑒𝑚𝑒𝑛𝑡𝑠 = 100 see that the lost mass is actually due to the creation of a gas called carbon dioxide. EXAMPLE with FORMULA Joseph-Louis Proust (1754 – 1826) ​ French Chemist who proved that the relative quantities of 1.​ A carbon dioxide which composed of one atom of carbon and two atoms of oxygen. A chemist measured the mass of each any given pure chemical compound’s constituent element present in 10 grams of CO2. Calculate the percentage elements remain invariant, regardless of the compound’s composition of CO2 by mass if the carbon is 2.73g and oxygen source. is 7.27g. ​ In his law, it states that in a given compound, the SOLUTION: elemental components are in a fixed ratio. 2.73𝑔 % 𝑜𝑓 𝑐𝑎𝑟𝑏𝑜𝑛 = 10𝑔 𝑥 100 = 27. 3% 7.27𝑔 Law of Definite Proportion % 𝑜𝑓 𝑜𝑥𝑦𝑔𝑒𝑛 = 10𝑔 𝑥 100 = 72. 7% ​ Also known as constant proportion. This law states that “ 27. 3% 𝑜𝑓 𝑐𝑎𝑟𝑏𝑜𝑛 + 72. 7% 𝑜𝑓 𝑒𝑥𝑦𝑔𝑒𝑛 = 100% 𝑜𝑓 𝑐𝑎𝑟𝑏𝑜𝑛 𝑑𝑖𝑜𝑥𝑖𝑑𝑒 Different samples of any pure compound contain the same elements in the same proportions by mass”. ​ Given compound always contains same proportions of Law of Multiple Proportion elements by mass. ​ The masses of one element combined with a fixed mass of another element in a different compound of the two Copper (II) Carbonate elements are in the ration of small whole numbers. 103g CuCO3 → 53g Cu+ 40g O2 + 10g C Copper (II) carbonate → Copper ​ When two elements form a series of compounds, the + Oxygen + Carbon ratios of the masses of the second element that combine ​ The elemental composition of a pure substance and the with one gram of the first element can always be reduced relative amounts of each element in a compound never to small whole numbers. varies or doesn’t vary. ​ If the two elements, A and B form more than one EXAMPLES compound, the masses of B that combine with a given mass of A are in the ratio of small whole numbers. 1.​ Consider the elements carbon and oxygen. Combined in one way, they form the familiar compound carbon dioxide. In every sample of carbon dioxide, there are 32.0g of oxygen EXAMPLES present for every 12.0g of carbon. By dividing 32.0 by 12.0 , this simplifies to a mass ratio of oxygen to carbon of 2.66 to 1. There is another compound that forms from the 1.​ Pure water will always contain hydrogen and oxygen in a fixed combination of carbon and oxygen called carbon monoxide. mass ratio (a gram of water consists of approximately 0.11 Every sample of carbon monoxide contains 16.0g of oxygen grams of hydrogen and 0.88 grams of oxygen, the ratio is 1:8). for every 12.0g of carbon. This is a mass ratio of oxygen to General Chemistry 1: MIDTERM NOTRE DAME OF GREATER MANILA — ACADEMIC SCHOOL YEAR 2024-2025 10  GENERAL CHEMISTRY 1 CHEM1-MID FIRST SEMESTER | ACADEMIC YEAR 2024-2025 | MR. MARK BENEDICK TOPINIO carbon of 1.33 to 1. In the carbon dioxide, there is exactly Subatomic Particles twice as much oxygen present as there is in the carbon ​ Sub-atomic particles are any of various self- contained monoxide. units of matter or energy that are the fundamental constituents of all matter. Figure 1: Parts of an atomic structure 2.​ In carbon monoxide, on the left, there is 1.333g of oxygen for every 1g of carbon. In carbon dioxide, on the right, there is 2.666g of oxygen for every gram of carbon. So the ratio of oxygen in the two compounds is 1:2, a small whole number ratio. EXAMPLE with SOLUTION 1.​ Ferrous chloride (FeCl2 ), 56 amu of iron is combined with 72 amu of chlorine, while in Ferric chloride (FeCl3 ), 56 amu of iron is combined with 105 amu of chlorine. SOLUTION: 𝐹𝑒𝐶𝑙2 → 56 𝑎𝑚𝑢 𝑜𝑓 𝐼𝑟𝑜𝑛 + 72 𝑎𝑚𝑢 𝑜𝑓 𝐶ℎ𝑙𝑜𝑟𝑖𝑛𝑒 𝐹𝑒𝐶𝑙2 → 56 𝑎𝑚𝑢 𝑜𝑓 𝐼𝑟𝑜𝑛 + 105 𝑎𝑚𝑢 𝑜𝑓 𝐶ℎ𝑙𝑜𝑟𝑖𝑛𝑒 𝐹𝑒𝐶𝑙2 56 𝑎𝑚𝑢 𝐼𝑟𝑜𝑛 = 𝐹𝑒𝐶𝑙3 56 𝑎𝑚𝑢 = 1: 1 Figure 2: Parts of an atomic symbol 𝐹𝑒𝐶𝑙2 72 𝑎𝑚𝑢 𝐶ℎ𝑙𝑜𝑟𝑖𝑛𝑒 = 𝐹𝑒𝐶𝑙3 105 𝑎𝑚𝑢 = 2: 3 − Therefore, with the fixed mass of iron, chlorine has a ratio of 2:3. Electrons (𝑒 ) ​ Discovered by Joseph John Thomson in the early 1890s. ​ Discovered using the Cathode Ray Tube Experiment. −19 −28 ​ Charge =− 1. 6022 𝑥 10 C, Mass = 9. 11𝑥10 𝑔 Atomic Structure & Subatomic Particles ​ Located at a specific energy level outside the nucleus. Atoms + Protons (𝑝 ) ​ The fundamental unit of matter or the building blocks of matter. ​ Discovered by Ernest Rutherford in the early 1900s. ​ It is made up of smaller particles which referred to as ​ Discovered using the Gold Foil Experiment. sub-atomic particles. ​ Located inside the nucleus. ​ It is the smallest particle of an element that can exist 𝑜 either alone or in combination with other atoms of the Neutrons (𝑛 ) same or of another element. ​ Discovered by James Chadwick. ​ Discovered using the Beryllium Foil Experiment. General Chemistry 1: MIDTERM NOTRE DAME OF GREATER MANILA — ACADEMIC SCHOOL YEAR 2024-2025 11  GENERAL CHEMISTRY 1 CHEM1-MID FIRST SEMESTER | ACADEMIC YEAR 2024-2025 | MR. MARK BENEDICK TOPINIO ​ Located inside the nucleus. ​ Atoms with the same number of protons but with different number of neutrons. Chemical Notation of Atoms ​ It can also be an atom with the same atomic number but ​ When an atom GAINS one or more electrons, it becomes different mass number. NEGATIVELY charged because it now holds more electrons ​ Isotopes of a given element has the same chemical than Protons. properties due to their same atomic number, but they ​ When an atom LOSES one or more electrons, it becomes differ in terms of their stabilities. POSITIVELY charged because it now holds fewer electrons ​ Some isotopes are unstable leading them to release than Protons. energy spontaneously through the process of decay. This is the reason why most of the isotopes are radioactive. ​ The only exception is the hydrogen where all of its isotopes has its name. Atomic Average Mass ​ The sum of the masses of the isotopes of a given element, each multiplied by its natural abundance. ​ The unit is amu (atomic mass unit). Figure 3: Chemical Notation of Atoms FORMULA 1.​ Key Terms ❖​ Atomic Number – Refers how many protons an atom of an element has. ❖​ Atomic Mass – Refers to the average mass of all the isotopes of an element. ❖​ Mass Number – The sum of protons and EXAMPLE with FORMULA neutrons in an atom. ❖​ Atomic Symbol – All elements have their own 1.​ Carbon-12 makes up 98.89% of naturally-occurring carbon. symbol, and it consist of a single capital letter Carbon-13 makes up 1.11% of naturally occurring carbon. Use followed by lower case letter/s. this information to determine the average atomic mass of carbon. 2.​ Key Concepts SOLUTION: ❖​ Atomic Number = Number of Protons Using the formula provided above… ❖​ Number of Electrons = Number of Protons 𝑎𝑚𝑢 = 12 𝑎𝑚𝑢 0. 9889 + (13𝑎𝑚𝑢)(0. 0111) ❖​ Proton — Electrons = Net Charge 𝑎𝑚𝑢 = 11. 8668 𝑎𝑚𝑢 + 0. 1443 𝑎𝑚𝑢 ❖​ Mass – Protons = Number of Neutrons 𝑎𝑚𝑢 = 12. 0111 𝑎𝑚𝑢 ❖​ Number of Protons + Number of Neutrons = Mass Number THEREFORE: ❖​ Nuclide – It refers to any atomic species 𝑎𝑚𝑢 = 12. 0111 𝑎𝑚𝑢 characterized by the number of protons and 2.​ Elemental bromine is made up of two naturally the number of neutrons. occurring isotopes Br-79 and Br-81 which have an atomic mass of 78.918 and 80.916 amu., respectively. The average atomic mass of Bromine is 79.9. Calculate KEY FORMULAS the percent abundance of each isotope. SOLUTION: Mass number = atomic Number of protons = Number Using the formula provided above… number + number of neutrons of electrons 79. 9 𝑎𝑚𝑢 = 78. 918 𝑎𝑚𝑢 𝑥 + (80. 916 𝑎𝑚𝑢)(1 − 𝑥) Number of neutrons = mass Net charge = number of 79. 9 amu = 78. 918 amu x + 80. 916 − 80. 916 amu x number - atomic number protons - number of electrons 79. 9 𝑎𝑚𝑢 − 80. 916 𝑎𝑚𝑢 = 78. 918 𝑎𝑚𝑢𝑥 − 80. 916𝑥 − 1. 016 𝑎𝑚𝑢 = − 1. 998 𝑎𝑚𝑢 𝑥 0. 509 = 𝑥 Isotopes General Chemistry 1: MIDTERM NOTRE DAME OF GREATER MANILA — ACADEMIC SCHOOL YEAR 2024-2025 12  GENERAL CHEMISTRY 1 CHEM1-MID FIRST SEMESTER | ACADEMIC YEAR 2024-2025 | MR. MARK BENEDICK TOPINIO 𝑥 = 0. 509 𝑥 100 = 50. 9% 1 − 𝑥 = 1 − 0. 509 0. 491 𝑥 100 = 49. 1% THEREFORE: 𝐵𝑟 − 79 = 50. 9% 𝐵𝑟 − 81 = 49. 1% Molecules and Ions Figure 4: Parts of chemical formula Molecules ​ A molecule is an aggregate of two or more atoms in a Binary Ionic Compounds definite arrangement held together by chemical bonds. ​ These are only made up of a metal (cation) and a ​ It is formed by chemical combination of atoms. non-metal (anion). ​ It is the smallest unit of a substance and shows all the ​ Retain the name of the 1 st element then remove the properties of that substance. ending of the second element and add “ -ide”. Classifications EXAMPLES 1.​ Monoatomic Molecule (Mono = one) ​ A molecule consist of one atom. MgO ​ Mg = Magnesium (retain its 2.​ Diatomic Molecule (Di = two) (Magnesium Oxide) name) ​ O = Oxygen (Change the ending “ ​ A molecule consist of two atoms. -ide” to Oxide 3.​ Triatomic Molecule (Tri = three) ​ A molecule consist of three atoms. Li3N ​ Li = Lithium (retain its name) 4.​ Polyatomic Molecule (Poly = many) (Lithium nitride) ​ N = Nitrogen (Change the ending “ -ide” to Nitride ​ A molecule consist of many atoms. NaCl ​ Na = Sodium (retain its name) Ions (Sodium Chloride) ​ Cl = Chloride (Change the ending “ -ide” to Chloride) ​ An ion is an atom, or group of atoms, that has a net positive or negative charge. ​ Ca = Calcium(retain its name) CaCl 2 ​ These are atoms with a charge. (Calcium Chloride) ​ Cl = Chloride (Change the ending ​ These are species carrying either a positive or negative “ -ide” to Chloride) charge. ​ Ions are formed by the addition or removal of electron in ​ In writing chemical formula of an ionic compound, the an atom. charges are very important because the numerical charge of one element would be the subscript of the other. Types ➔​ To identify the chemical formula of a chemical compound, 1.​ Cation (Metals) – Positive ion formed by the removal of 1.​ we need to know first the symbol. electron. 2.​ we need to know the charge of ions present in 2.​ Anion (Nonmetals) – Negative ion formed by the addition this compound of electron. 3.​ then do the CRISS-CROSS METHOD. Naming Chemical Compounds NOTE: When writing a chemical formula, always remember that we can reduce the subscripts to its lowest term. If that is the Chemical Formula case, always consider your anion because it will be an indicator if ​ A way of representing the composition of a molecule or the subscripts were reduced. an ion. ​ It is consisting of the symbols of the atoms that makes up the molecules. EXAMPLES ​ If there is more than one atom present, a numerical SUBSCRIPT is used. General Chemistry 1: MIDTERM NOTRE DAME OF GREATER MANILA — ACADEMIC SCHOOL YEAR 2024-2025 13  GENERAL CHEMISTRY 1 CHEM1-MID FIRST SEMESTER | ACADEMIC YEAR 2024-2025 | MR. MARK BENEDICK TOPINIO Figure 5: List of the most common polyatomic ions EXAMPLES Ionic Compounds with Polyatomic Anions ​ Naming ionic compounds with monoatomic cation & polyatomic anion. ​ Retain the name of both ions. ​ The way of writing chemical formula for an ionic compound with polyatomic ion is just the same with binary compounds. ​ The only difference is that we should be mindful on the use of parenthesis. General Chemistry 1: MIDTERM NOTRE DAME OF GREATER MANILA — ACADEMIC SCHOOL YEAR 2024-2025 14  GENERAL CHEMISTRY 1 CHEM1-MID FIRST SEMESTER | ACADEMIC YEAR 2024-2025 | MR. MARK BENEDICK TOPINIO Ionic Compounds with Variable Valence ​ Naming with variable valence is named using CLASSICAL and STOCK NAMING SYSTEM. 1.​ Classical Naming System ➔​ It uses the Latin root name of the cation changing its ending to “ -ic” (with greater charge) and “ -ous” (with lesser charge). ➔​ The anion will be named the same way. For 2.​ Stock of Naming System monoatomic, change the ending to “ - ide” and ➔​ It uses the Original name of cation and roman for polyatomic, retain its name. numeral indicates the charge before the non-metal. ➔​ The anion will be named the same way. For monoatomic, change the ending to “ - ide” and for polyatomic, retain its name. EXAMPLES Figure 6: List of common classical names EXAMPLES Binary Covalent Compounds ​ Binary covalent compounds are made up of two non-metals. ​ RULE: 1.​ Always have a suffix of “ -ide” for the second non-metal. 2.​ Use prefixes to indicate the number of each element present in the compound. 3.​ DO NOT use the prefix “mono- ” in the first non-metal if it only has a subscript of 1. General Chemistry 1: MIDTERM NOTRE DAME OF GREATER MANILA — ACADEMIC SCHOOL YEAR 2024-2025 15  GENERAL CHEMISTRY 1 CHEM1-MID FIRST SEMESTER | ACADEMIC YEAR 2024-2025 | MR. MARK BENEDICK TOPINIO Rules in Naming Acids without Oxygen 1.​ Write hydro as the first name of the acid to represent the hydrogen atom. 2.​ For the anion, remove the suffix “ -ide” and replace it with “ -ic”. EXAMPLES Figure 7: List of common numer to greek prefix EXAMPLES Rules in Naming Acids with Oxygen 1.​ Do not use Hydro. 2.​ If the suffix of a polyatomic anion is “ -ate”, change it to “ -ic”. 3.​ If the suffix of a polyatomic anion is “ -ite”, change it to “ -ous”. EXAMPLES Naming Acids ​ Acids are compound in which one or more H+ are bonded to a negative ion. It is considered as covalent due to the presence of non-metals. General Chemistry 1: MIDTERM NOTRE DAME OF GREATER MANILA — ACADEMIC SCHOOL YEAR 2024-2025 16  GENERAL CHEMISTRY 1 CHEM1-MID FIRST SEMESTER | ACADEMIC YEAR 2024-2025 | MR. MARK BENEDICK TOPINIO Supplemental Videos and Websites ​ How To Name Ionic Compounds With Transition M… Steps in Balancing Chemical Equations ​ How To Name Ionic Compounds In Chemistry 1.​ Take note of the elements present in the reactant and ​ Naming Ionic Compounds product side. ​ Naming ions and ionic compounds (video) | Khan 2.​ Count the no. of atoms each element present in reactant Academy & product side. ​ 5.8: Naming Molecular Compounds - Chemistry LibreTexts 3.​ Add coefficients to make the number of atoms equal and ​ 5.7: Naming Ionic Compounds - Chemistry LibreTexts double check if balance. Balancing Chemical Equation EXAMPLES Stoichiometry ➔​ “Stoicheion” means element and “metron” means measure. ➔​ Section of chemistry that involves using relationships between reactants and/or products in a chemical reaction to determine desired quantitative data. ➔​ The calculations are based on balanced equation. ➔​ Amount of Reactants and Products in Reactions. ➔​ The study of quantitative relationships between the number of REACTANTS used and PRODUCTS formed in a chemical reaction. Chemical Equations ​ An equation that uses chemical symbols and formulas to represent a chemical reaction. ​ A shorthand way of writing a chemical reaction using chemical symbols and formulas. Figure 8: Sample of chemical equation Symbols in Chemical Equations General Chemistry 1: MIDTERM NOTRE DAME OF GREATER MANILA — ACADEMIC SCHOOL YEAR 2024-2025 17  GENERAL CHEMISTRY 1 CHEM1-MID FIRST SEMESTER | ACADEMIC YEAR 2024-2025 | MR. MARK BENEDICK TOPINIO Molar Mass ​ It is also known as molecular mass/molecular weight; this is the mass in grams of 1 mol of the substance. ​ For compounds, molar masses are calculated by summing the atomic masses of all atoms appearing in a chemical formula. KEY FORMULAS Number of 𝑚𝑎𝑠𝑠 moles 𝑛= 𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 Mass 𝑚𝑎𝑠𝑠 = 𝑛𝑜 𝑜𝑓 𝑚𝑜𝑙𝑒𝑠 * 𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 Moles to no. of entities TIPS: No. of 1.​ Balance first the Carbon. entities to 2.​ Leave Hydrogen second to the Last by getting the half of mass the number of hydrogen. 3.​ Balance the number of Oxygen for Last. EXAMPLES The Mole Concept Number ​ Mole is the amount of substance containing the same of moles number of elemental entities (atoms, molecules, ions, etc.) as the number of atoms in a sample of pure 12C weighing exactly 12 g. ​ It is abbreviated as mol, the SI unit for the amount of chemical species. ​ It is always associated with chemical formula and represents Avogadro’s number which is 6.022 x 1023 of particles. ​ 1 mole = 6.022 x 1023 particles/entities. General Chemistry 1: MIDTERM NOTRE DAME OF GREATER MANILA — ACADEMIC SCHOOL YEAR 2024-2025 18  GENERAL CHEMISTRY 1 CHEM1-MID FIRST SEMESTER | ACADEMIC YEAR 2024-2025 | MR. MARK BENEDICK TOPINIO ❖​ The compounds will always add up to 100%, so in a binary Mass compound, you can find the % of the first element, then do 100%-(% first element) to get (% second element) ❖​ If using a calculator, you can store the overall molar mass to a variable such as "A". This will speed up calculations, and reduce typographical errors. Example With The Use Of Steps Above Moles to 1.​ In mammals, lactose(milk sugar) is metabolized to no. of glucose (C6H12O6 ) ,the key nutrient for generating entities chemical potential energy. What is the mass percent of each element in glucose? SOLUTION: 1.​ Find the molar mass of all the elements 𝐶 = 12. 012𝑔 𝑥 6 = 72. 072𝑔 𝐻 = 1. 008𝑔 𝑥 12 = 12. 10𝑔 No. of entities to 𝑂 = 16𝑔 𝑥 6 = 96𝑔 mass 2.​ Find the molecular mass of the entire compound. C6H12O = 72. 072𝑔 + 12. 10𝑔 + 96𝑔 = 180. 172𝑔/𝑚𝑜𝑙 3.​ Divide the component's molar mass by the entire molecular mass 72.072𝑔 𝐶 = 180.172 𝑔/𝑚𝑜𝑙 × 100 = 40. 002% Percent Composition 12.10𝑔 𝐻 = 180.172 𝑔/𝑚𝑜𝑙 × 100 = 6. 72% ​ Mass percent can also be used to calculate the mass of a 96 𝑔 particular element in any mass of a compound. 𝑂 = 180.172 𝑔/𝑚𝑜𝑙 × 100 = 53. 28% ​ Percent of all elements in a compound must equal to 100. ​ The percent by mass of each element in a compound is 4.​ Add all computed molar mass. If the result is near called the percent composition of a compound. 100, the computation is correct. 40. 002% + 6. 72% + 53. 28% = 100. 002% THEREFORE: Glucose is 40.002% C, 6.72% H, and 53.28% O by mass. Figure 9: Mass of Element and percent by mass formula Steps to Solve Empirical and Molecular Formula 1.​ Find the molar mass of all the elements in the compound 1.​ Empirical Formula - elements are expressed in simplest in grams per mole. whole- number ratio of atoms in a given compound. 2.​ Find the molecular mass of the entire compound. Example: For Acetylene the empirical formula is CH 3.​ Divide the component's molar mass by the entire 2.​ Molecular Formula - gives the real number of atoms molecular mass. present in the compound. Example: For Acetylene the 4.​ You will now have a number between 0 and 1. Multiply it empirical formula is C2H2. by 100 to get percent composition! 5.​ Add all computed molar mass. If the result is near 100, the Steps to Solve computation is correct. 1.​ Assume a 100g sample of the compound so that the given percentages can be directly converted into grams. Tips to Solve 2.​ Use each element's molar mass to convert the grams of each element to moles. General Chemistry 1: MIDTERM NOTRE DAME OF GREATER MANILA — ACADEMIC SCHOOL YEAR 2024-2025 19  GENERAL CHEMISTRY 1 CHEM1-MID FIRST SEMESTER | ACADEMIC YEAR 2024-2025 | MR. MARK BENEDICK TOPINIO 3.​ In order to find a whole-number ratio, divide the moles of