Midterm Chemistry 1, 2, 3, 4, 5, 6, 7 PDF

MODULE 1 CHEMISTRY is the study of matter and the changes it undergoes, it is divided in different branches, including organic, inorganic, physical, analytical, biochemistry. MATTER is anything that takes up space. MASS is the measurement that reflects the amount of matter. SCIENCE is the use of...

MODULE 1 CHEMISTRY is the study of matter and the changes it undergoes, it is divided in different branches, including organic, inorganic, physical, analytical, biochemistry. MATTER is anything that takes up space. MASS is the measurement that reflects the amount of matter. SCIENCE is the use of evidence to construct testable explanations of natural phenomena An HYPOTHESIS is a testable expectation and prediction of a phenomena. A THEORY is an explanation of a phenomena based on observations over time. It can be modified. SCIENTIFIC LAW is a relationship in nature supported by human experiments PURE RESEARCHis for scientists to gain knowledge. APPLIED RESEARCH is for scientists to solve a specific problem. A MODEL is a visual, verbal explanation of data. UNITS BASE UNITS are a system of measurements. DERIVED UNIT is a unit derived from the combination of two units. Example is meter per seconds. WEIGHT= measure of the amount of matter and the gravitational force exerted on a object, TIME= second LENGTH= meters MASS= KILOGRAMS TEMPERATURE= kelvin Fahrenheit is used in the usa water freeze at 32 degrees f and boils at 212 degrees f Celsius is used in the rest of the world freezing point of water is 0 boiling point is 100 Kelvin is used by scientists water freezes at 273.15 and boil at 373.15 VOLUME is the space occupied by an object. Measure is liters DENSITY is a physical property defined as the amount of mass per unit volume Formula is mass/ volume Module 2 Physical properties: color, odor, density, mass, solubility, state, melting point, boiling point Chemical properties: reactivity, flammability, acidity/basicity, chemical composition. In essence, physical properties can be observed without altering the substance, while chemical properties are concerned with how the substance changes in the context of chemical reactions. STATES OF MATTER: SOLid: has both shape and volume. Particles in a solid are tightly packed and a solid does not take space in a container. LIQUID: no shape but has volume. It takes the space of a container and particles move past each other. GAS: neither shape of fixed volume. It takes up the space and volume of the containers, particles in gas are very far apart. PLASMA: it is a form of matter that results from the particles of a gas becoming ionized and broken apart into smaller particles. Plasma is an ionized gas. PHYSICAL CHANGE alters the substance without changing its composition. For example crumpling aluminium foil into a ball doesn't change the composition, it is still aluminium A phase change is the change from one state of matter to the other. CHEMICAL CHANGE OR CHEMICAL REACTION is the changing of one or more substances into a new one. DIFFERENCE IS THAT IN THE PHYSICAL CHANGE THERE IS NOT CHANGE IN THE COMPOSITION WHILE IN A CHEMICAL CHANGE YES. STARTING SUBSTANCES ARE CALLED reactants WHIòE THE RESULTS ARE CALLED products LAW OF CONSERVATION OF MASS mass is neither created nor destroyed, it is conserved. The products equal to the reactants AN ELEMENT is a pure substance that cannot be simplified into a simple substance in a physical or chemical change. THE PERIODIC TABLE the first version was created by mendeleev. Nowadays it is organized in horizontal rows called period and vertical columns called groups. COMPOUNDS are made up of two or more elements combined chemically in a fixed ratio. Example is h20. They can be broken down into smaller substances by chemical means. Separating a compound into its element requires heat or electricity. Two elements combined together are different compared to them separately. THE LAW OF DEFINITE PROPORTION says that a compound has the same elements in the same proportions by mass, no matter how large or small the sample. The mass of the compound is equal to the sum of the masses of the elements that make up the compound. THE PERCENT BY MASS is the ratio of the mass of each element to the total mass of the compound expressed in percent. Percent by mass= mass of the element/ mass of the compounds x 100 THE LAW OF MULTIPLE PROPORTIONS states that when elements combine to form compounds, the results will depend on the proportions (ratios) MIXTURES when two or more substance combine HETEROGENEOUS MIXTURE substances remain distinct and aren't uniform HOMOGENEOUS MIXTURE subattributes are uniform and are also referred as solutions SEPARATION TECHNIQUES FILTRATION is a separation technique where a porous barrier is used to separate solids from liquids. Heterogeneous mixtures. EXAMPLE is a mixture of sand and water. DISTILLATION is based on different boiling points. Homogenous mixtures. Example distillation of water  ​ First, the liquid heats up and turns into a gas ​ Then, the gas cools down and turns back into a liquid ​ So, distillation is the process of going from liquid → gas → liquid again. Sublimation is when a substance goes directly from a solid to a gas without turning into a liquid first. Based on pressure and temperature EXAMPLE DRY ICE CHROMATOGRAPHY works by passing the mixture through a material that allows some substances to move faster than others, helping to separate them. For example, it's like how a drop of ink spreads out into different colors on a piece of paper! Based on solubility. CRYSTALLIZATION Crystallization is a way to separate a solid from a liquid. The solution needs to dissolve, for example sugar and water. Then the liquid needs to dry in order to make the solòid form crystals. Module 3 ARISTOTLE believed that the atom didn’t exist and that empty space didn't exist. He proposed that all matter was made up of four fundamental elements: earth, water, air, and fire. DEMOCRITUS instead believed that matter was made up of tiny particles called atoms. DALTON’S ATOMIC THEORY ​ Matter is composed of tiny small particles called atoms ​ Atoms and invisible and indestructible ​ Atoms of a given elements are the same in size, mass, properties ​ Atoms of a specific element are different from those of another elements ​ Atoms combine to form compounds through ratios ​ In chemical reactions, atoms are separated, combined or rearranged WRONG POSTULATES OF DALTON’S THEORY Atoms are indivisible and indestructible: ​ Dalton stated that atoms cannot be divided into smaller parts. ​ This is incorrect because atoms are made up of subatomic particles (protons, neutrons, and electrons). Additionally, nuclear reactions show that atoms can be split. Atoms of a given element are identical in mass and properties: ​ Dalton believed all atoms of the same element are identical. ​ This is incorrect because isotopes of elements exist, where atoms of the same element have different masses (due to a different number of neutrons). Cathode rays are streams of negatively charged particles that are emitted from the cathode in a vacuum tube. These rays were later identified as electrons, meaning cathode rays and electrons are the same. This is because experiments, like those conducted by J.J. Thomson showed that cathode rays consist of particles much smaller than atoms with a negative charge, and these particles are what we now know as electrons. Plum Pudding Model (J.J. Thomson, 1904): ​ Thomson proposed that the atom is a sphere of positive charge with negatively charged electrons scattered within it, like "plums in a pudding." ​ This model did not include the nucleus Rutherford's Gold Foil Experiment (1909): ​ Conducted by Ernest Rutherford, Hans Geiger, and Ernest Marsden. ​ Alpha particles were directed at a thin gold foil. Most passed through, but some were deflected at large angles, and a few bounced straight back. ​ Revealed the presence of a nucleus at the center of the atom Rutherford's Nuclear Model (1911): ​ Based on the experiment, Rutherford concluded that: 1.​ Atoms have a dense, positively charged nucleus where most of the mass is concentrated. 2.​ Electrons orbit this nucleus, and most of the atom is empty space. ​ This model replaced the Plum Pudding Model. Alpha particles contain two protons and two neutrons, and has 2 plus change, which explains why they are attracted to the negatively charged plate. The element associated with alpha particles is helium (He). Alpha particles are essentially helium nuclei, consisting of two protons and two neutrons, but without electrons. Helium atom: neutral; consists of two protons, two neutrons, and two electrons. They are emitted during the radioactive decay of heavy elements like uranium and radium.  -​ Alpha particles are useful in various applications, such as in smoke detectors, where they help ionize air to detect smoke particles. -​ They are also used in radiation therapy to target and destroy cancer cells. Alpha particles are slower and heavier compared to beta particles and gamma rays. They have a mass roughly equivalent to four protons. Due to their strong ionizing capability, alpha particles are unable to penetrate deeply into matter and can be blocked by thin layers of skin or even air. On the other hand, beta particles are fast-moving, high-energy electrons much lighter than alpha particles and carrying either a positive or negative charge. Beta radiation can cause harm, such as skin burns, as it can pass through several millimeters of tissue. In contrast, gamma rays are emitted by the most energetic and hottest objects in the universe. Their highly penetrating nature comes from their extremely small photons, which allow them to pass through matter. Gamma rays are the most hazardous form of radiation because they can penetrate directly through the skin without causing ionization. However, when ingested or inhaled, alpha particles become more harmful than gamma rays. Z is the representation of the atomic number (number of protons) The number of electrons are equal to the number of protons Atomic mas= z + electrons+ neutrons Mass number= z + neutrons (mass of the nucleus) 17. Boron (B) has two naturally occurring isotopes: boron-10 (abundance = 19.8%, mass= 10.013 amu) and boron-11 (abundance = 80.2%, mass = 11.009 amu). Calculate the atomic mass of boron. Fist calculate the mass contribution= atomic mass x percent abundance Mass contribution for boron-10= 10.013 amu x 19.8%= 1,982 Mass contribution for boron-11= 11.009 amu x 80.2%= 8,829 Second, calculate the atomic mass= mass contribution 1 + mass contribution 2 Average atomic mass of boron= 1,982 + 8,829= 10,811 amu 18. CHALLENGE Nitrogen has two naturally occurring isotopes, N-14 and N-15. Its atomic mass is 14.007. Which isotope is more abundant? Explain your answer. N14 has 99.64% abundance percentage and N15 has 0.35% abundance percentage. Radioactive Decay: Radioactive decay is the process by which an unstable atomic nucleus loses energy by emitting radiation. This happens because the nucleus is trying to become more stable. During this process, the nucleus releases particles or energy in the form of alpha particles, beta particles, or gamma rays. Alpha Particles: Alpha particles are made of 2 protons and 2 neutrons (like a helium nucleus). They have a positive charge, are heavy, and move slowly. Because of their size, they can be stopped by a sheet of paper or even skin. Charge is of +2 Beta Particles: Beta particles are high-energy electrons or positrons. They have a negative or positive charge and are much lighter and faster than alpha particles. They can penetrate further, needing materials like aluminum to stop them. Gamma Rays: Gamma rays are not particles but high-energy electromagnetic waves. They have no mass and no charge, making them very penetrating and dangerous. They need heavy shielding, like lead or thick concrete, to block them. Module 4 Lesson 1 Electromagnetic radiation= a form of energy that exhibits a wavelike behavior as it travels through space. Wavelength: The distance between two identical points (like two peaks) in a wave. Frequency: How many times a wave repeats (or cycles) in one second. Amplitude: The height of a wave from its middle (rest position) to its peak (or trough). It shows how strong or intense the wave is. As frequency increases, wavelength decreases and viceversa. Fission: splitting a large atomic nucleus into two smaller ones and releasing more neutrons. These neutrons then hit other atoms, causing a *chain reaction*.When the atom splits, a small amount of its mass turns into energy. If too many atoms split too quickly, it can cause an explosion, like a nuclear bomb. Fusion: Combining two small atomic nuclei to form a larger one, also releasing a lot of energy. It’s like merging two small drops of water into one big drop. Fission powers nuclear reactors, while fusion powers the sun! Speed of light or c= 3.00 x 108 m/s c= λ x f continuous spectrum has no interruptions and contains all wavelengths of light without any gaps, appearing as a smooth blend of colors. White light produces a continuous spectrum. frequency and wavelength are opposites: when one goes up, the other goes down. The faster the waves pass by (high frequency), the more energy they carry. Violet light has fast waves (high frequency) and carries more energy than red light, which has slow waves (low frequency). Lastly, all these waves, whether they’re light, radio signals, or X-rays, are part of this giant collection of waves called the electromagnetic spectrum. Atomic emission spectrum= a series of lines of individual colors. It’s created when atoms or gases are excited and produce light at certain wavelengths. Quantum = minimum amount of energy that can be gained or lost by an atom E of a quantum= h x f h= planck’s constant, which has a value of 6.626 x 10-34 J Photoelectric effect= when light of a certain frequency strikes a metal surface and ejects electrons, or photoelectrons. Einstein, to explain the photoelectric effect, proposed -that light has a dual nature. A beam of light has both wavelike and particle-like properties. Photon= a massless particle that carries quantum energy. A photon’s energy depends on its frequency. The energy of a photon just has a certain threshold value to cause the ejection of a photoelectron from the surface of a metal. Lesson 2 Ground state= an atom being at its lowest energy level. When the atom gains energy, or is in an excited state, the electron moves to a higher state Quantum number= n, is assigned to each orbit. Bohr says that when a hydrogen atom is in the ground state when its single electron is in the n= 1 orbit, also called the first energy level. Electrons move in circular paths called orbits or shells. Bohr suggested that if you know which orbit an electron is in, you can determine its distance from the nucleus and its general path, like a planet orbiting the sun. When the atom is in an excited state, the electron can drop from the high energy level to a lower one, and a photon is emitted. The photons will have different wavelengths and frequencies; this makes photons of different energies produce different colors of light. The hydrogen atom’s electron can move only from one allowable orbit to another and can emit or absorb only certain amounts of energy. De Broglie was convinced that Bohr’s atomic model was incorrect, so he proposed another idea. He thought that tiny things like electrons could act like waves, just as light can (dual nature). De Broglie’s equation = h/ mv h = planck’s constant (6.626 x 10-34 J) m= mass of the particle v= velocity This means that if something is light and fast, its wavelength is big enough to notice. If it’s heavy and slow, the wavelength is too tiny. We can’t see the wave-like in big objects because their wavelengths are so tiny that they are impossible to measure. Heinsenberg discovered that measuring something small, like an electron, will always disturb it. This makes it impossible to know exactly where it is and how fast it’s moving at the same time. Heinsenberg’s uncertainty principle= when we try to measure the position where the particle is or its speed, we affect one or the other. You can’t get both with perfect accuracy. The only quantity of an electron’s orbit that can be determined is the probability of where it might be. Schrodinger’s model for the hydrogen atom seemed to apply equally well to atoms of other elements, unlike Bohr’s model. Quantum mechanical model of the atom= the atomic model in which electrons are treated like waves. Atomic orbital= the electron’s probable location.We don’t know exactly where an electron is, but we can say it’s somewhere in a three-dimensional area around the nucleus. Scientists use a math equation called a wave function to predict where an electron is most likely to be. The more dots or density in one area, the more likely the electron is in there. Scientists draw a boundary around 90% of the area where the electron is most likely to be. Principal quantum number= n tells about the main energy level where an electron can be found. As n increases, the size of the orbital gets larger, and the electron is further from the nucleus with more energy. The principal energy level contains a smaller energy division called energy sublevels. The number of energy sublevels in a principal energy level increases as n increases. Sublevels are labeled s,p,d or f according to the shape of the atom’s orbitals. Each orbital can hold a maximum of two electrons. s= -all spherical shapes (1s,2s,3s, etc) p= -dumbbell-shaped -Each principal energy level starting from n=2 has three p orbitals (2px , 2py, 2pz) -oriented along the x,y and z axes and all three p orbitals together can hold a total of 6 electrons. d= -complex shapes -They start appearing from the third principal energy level (n=3) and can be found in higher levels -Five distinct d orbitals with a unique shape and orientation: dxy, dyz, dzx, dx2-y2, dx2. -all the five orbitals together hold a total of 10 electrons. f= very complex shapes -​ Start appearing from the fourth principal energy level. -​ Have seven distinct orbitals. -​ Can hold up to 14 electrons across all f orbitals 1. Main Energy Levels The term “energy level” in physics and chemistry refers to the specific energies that electrons can have when occupying orbitals around an atom’s nucleus. These levels are often denoted by the principal quantum number (n), which describes the distance of an electron from the nucleus. The higher the energy level (or value of n), the further the electron is from the nucleus, and the more energy it possesses. 2. How Many Energy Levels Can There Be? elements have been observed to have energy levels up to n=7 in their ground state. 3. How Many Electrons in Each Energy Level? The maximum number of electrons in each energy level is given by the formula: 2n2 ​ n=1: 2 x 12= 4 electrons ​ n=2: 2 x 22= 8 electrons ​ n=3: 2 x 32= 18 electrons ​ n=4: 2 x 42= 32 electrons ​ n=5: 2 x 52= 50 electrons ​ n=6: 2 x 62= 72 electrons ​ n=7: 2 x 72= 98 electrons 4. What Energy Levels Can Include s and p Sublevels? The sublevels are denoted as s, p, d, and f. The availability of these sublevels in each energy level depends on the principal quantum number n: ​ n=1: only s, total of 2 electrons ​ n= 2: s and p, total of 8 electrons ( 2 from s, 6 from p) ​ n= 3: s, p and d, total of 18 electrons (2 from s, 6 from p, 10 from d) ​ n= 4 and higher: s,p,d, f holds up to 32 electrons the energy levels that can include s and p sublevels start from n=2 Lesson 3 Electron configuration= arrangement of electrons in an atom at its ground level. Less energy= more stable Aufbau principle= each electron occupies the lowest energy level available. Pauli exclusion principle= a maximum of two electrons can occupy each atomic orbital, only if electrons have opposite spins ↑ electron is spinning in one direction ↓ electron is spinning in the opposite direction ↑↓ filled orbital with a pair of electrons with opposite spins. Hund’s rule= single electrons with the same spin must occupy each equal energy orbital before electrons with opposite spins can occupy the same energy orbitals. Bromine 35 (Br) 1s2, 2s2, 2p6, 3s2, 3p6, 1s2, 2s2, 2px2, 2py2, 2pz2, 3s2, 3px2, 3py2, 3pz2, 3d10 4s2, 4px2, 4py2, 4pz1 n=1 only s sublevel n=2 s and p sublevels n=3 s, p, d sublevels n=4 s,p,d,f sublevels Strontium (Sr) 38 1s2 2s2, 2px2, 2py2, 2pz2 3s2, 3px2, 3py2, 3pz2, 3d10 4s2, 4p6, 5s2 n=5 is valence energy level 2 valence electrons 1s2,2s2,2p6, 3s2,3p6,4s2,3d10,4p6,5s2 Tungsten (W) 74 1s2 2s2, 2p6 3s2, 3p6, 3d10 4s2, 4p6, 4d10, 4f14 5s2, 5p6, 5d4 6s2 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6, 5s2, 4d10, 5p6, 6s2, 4f14, 5d4 Noble gas notation Aluminum 13 (Al) [Ne] 3s2, 3p1 Calcium 20 [Ar] 4s2  Chapter 5 Lesson 1 Old Periodic table was arranged by atomic mass, and elements with similar properties horizontally, Mendeleev proposed the periodic table arranged in order by increasing atomic number, Periodic law= there is a periodic repetition of physical and chemical properties when elements are arranged by increasing atomic number Columns are called groups: Atoms in the same group have similar chemical properties because they have the same number of valence electrons. There are 18 groups. Rows are called periods. There are 7 of them. Representative elements: group 1 and 2 and from 13 to 18. Transition elements: group 3 to 12. Alkali metals are group 1 Alkali earth metals group 2 Inner transition metals are lanthanide series and actinide series (in the bottom) Halogens group 17, very reactive Noble gases group 18, extremely unreactive. They are the most stable and don't do chemical reactions have already 8 electrons. Metals are good conductors of heat and electricity, malleable, and ductile, transition elements are metals. They lose electrons to become stable Nonmetals are bad conductors of electricity. Gain electrons to become stable. Metalloids have physicaò and chemical properties of both metals and nonmetals Lesson 2 Atoms in the same group have similar chemical properties because they have the same number of valence electrons. Group 13 to 18 have ten less valence electrons than their group number. S block: 1 and 2 groups P block: 13 to 18 D block: transition elements F block: inner transition elements Lesson 3 Atomic radius The distance from the nucleus of an atom to the outermost electron, representing the size of the atom. As you move left to right across a period, the atomic radius decreases because there is an increasing positive charge and the energy level remains the same As you go down a group, atomic radii increases because the energy level is higher and the size of the atom is bigger. ion= atom that gains or loses electrons to become stable Cation: positive ion (lost electron) METAL. The atom shrinks because the valence electrons are lost and the entire energy level as well. Anion: negative ion (gained electron) NONMETAL BUT NOT NOBLE GASES. Atom becomes bigger Trend of ionic radius: As you move left to right across a period, the size of the positive ions decreases. As you move down a group, there is an increase in ionic size Nonmetals tend to increase as you move from left to right across a period. Ionization energy energy required to remove an electron from a gaseous atom. First ionization energy is the energy required to remove the first outermost electron from an atom. Nonmetallic character increases as you go to the right, elements gain electrons, Ionization energy increases Trend of ionization energy Across a Period (Left to Right): Ionization energy increases because the number of protons in the nucleus increases, which strengthens the attraction between the nucleus and the electrons. This makes it harder to remove an electron as you move across a period. Down a Group: Ionization energy decreases because atoms have more electron shells as you go down a group. The outermost electrons are farther from the nucleus, and the shielding effect from inner electrons weakens the pull of the nucleus. This makes it easier to remove an electron Nonmetals gain electrons Octet rule atoms tend to gain, lose or share electrons in order to acquire a full set of eight valence electrons.. First period elements are an exception. Electronegativity is the ability of an element’s atoms to gain electrons. Decreases down a group because the nucleus is farther away from the atom , increases as you move left to right because they want to gain electrons.  Module 6 Lesson 1 Atoms combine to form compounds because they want to be stable. Stability happens when their outer layer of electrons (valence electrons) is “full.” This is why atoms either share, lose, or gain electrons. Ions are formed by the loss or gain of valence electrons.The number of prtomns remains unchanged during ion formation. Chemical bonds are the forces that hold two atoms together. They can share electrons (covalent bonds) or transfer electrons to form positive and negative ions (ionic bonds). Atoms with incomplete outer shells of electrons (valence electrons) bond to become stable, achieving a configuration like noble gases. Electron-dot structures are simple drawings that show an element’s valence electrons as dots around the element’s symbol. Example: Sodium (Na) has one dot, because it has one valence electron. When an atom loses one or more valence electrons, it forms a positive ion (cation). Example: Sodium (Na) has 1 valence electron. By losing it, sodium becomes Na⁺, which is positively charged. When an atom gains one or more electrons, it forms a negative ion (anion). Example: Chlorine (Cl) has 7 valence electrons. It gains 1 electron to form Cl⁻, becoming stable like argon. Metals are reactive because they lose their valence electrons easily. For example:Group 1 metals lose 1 electron to form ions like Na⁺ or K⁺Group 2 metals lose 2 electrons to form ions like Mg²⁺ or Ca²⁺. Transition metals can lose electrons from 4s and 3d orbitals. They form ions with different charges, like +2 or +3, because: ​ They lose electrons from their 4s and 3d orbitals. ​ The resulting ions are more stable due to the formation of half-filled or fully filled d-subshells. ​ The ability to achieve these stable electron configurations is a key reason why these metals exhibit multiple oxidation states. A pseudo-noble gas configuration happens when a transition metal loses its outer electrons (4s orbital) but still has a full d sublevel (like 3d¹⁰). This makes the ion stable, even though it doesn’t have a true noble gas configuration. Example: Zinc (Zn) ​ ​ Zinc: [Ar] 4s² 3d¹⁰. ​ ​ Zn²⁺ (loses 4s electrons): [Ar] 3d¹⁰. The full 3d sublevel makes Zn²⁺ stable. Nonmetals form ions by gaining electrons to fill their outer shell with 8 (a stable octet): Group Rules: ​ ​ Group 15: Gain 3 electrons → 3⁻ ion. ​ ​ Group 16: Gain 2 electrons → 2⁻ ion. ​ ​ Group 17: Gain 1 electron → 1⁻ ion. Lesson 2 In Ionic Bonds, Atoms transfer electrons to become stable.For Example: Sodium gives an electron to chlorine. Sodium becomes a positive ion, and chlorine becomes a negative ion. These compounds (like table salt, NaCl) are called ionic compounds. Binary Ionic Compounds are made of two elements: one metal (positive ion) and one nonmetal (negative ion). Example: Magnesium oxide (MgO). Ionic compounds are made of large numbers of positive and negative ions arranged in a specific ratio. This ratio depends on how many electrons are transferred between the metal and nonmetal atoms. The strong electrical attractions between positive and negative ions cause them to organize into a crystal lattice—a stable geometric arrangement ( a repetitive pattern). Each positive ion is surrounded by negative ions and vice versa, ensuring the system is neutral overall. The shape of the lattice can vary depending on the sizes and ratios of the ions. ​ Example: Sodium Chloride (NaCl): In NaCl, each sodium ion is surrounded by six chloride ions, and each chloride ion is surrounded by six sodium ions. This creates a repeating cubic pattern. Key Concept: The physical structure and shape of ionic crystals are determined by the ratio and arrangement of ions in the lattice. Ionic compounds exhibit several unique properties due to the strong ionic bonds in their lattice: High Melting and Boiling Points= Ionic bonds are very strong, requiring a lot of energy to break. As a result, ionic compounds have high melting points (e.g., NaCl melts at 801°C) and boiling points (over 1400°C in many cases). Hardness= Ionic compounds are hard and rigid because the ions are held in fixed positions within the lattice. Colors= Many ionic compounds are vividly colored, especially those containing transition metals. These colors result from the interaction of light with the electrons in the metal ions. The ability of ionic compounds to conduct electricity depends on whether the ions can move freely. Solid State: Ions are locked in place within the lattice, so ionic solids do not conduct electricity. Liquid State or in Solution: When melted or dissolved in water (so in aqueous states), the ions are free to move, allowing the compound to conduct electricity. Such substances are called electrolytes. Ionic bond formation lowers the system's energy and releases heat, making the process exothermic. Lattice energy is the energy required to separate one mole of ions in an ionic compound. It reflects the strength of the bonds in the crystal lattice: ​ Ion Size: Smaller ions can pack more closely, increasing the force of attraction and lattice energy. For instance, a lithium ion (Li⁺) has a stronger attraction than a larger potassium ion (K⁺), leading to higher lattice energy. ​ Ion Charge: Higher charges on ions result in stronger attractions and greater lattice energy. For example, magnesium oxide (MgO) has much higher lattice energy than sodium fluoride (NaF) because the charges on Mg²⁺ and O²⁻ are greater than those on Na⁺ and F⁻. Key Concept: Lattice energy depends on ion size and charge—the smaller and more charged the ions, the stronger the ionic bond. Lesson 3 Formula Units ​ Ionic compounds are made of positive (cation) and negative (anion) ions that stick together in a pattern. ​ The chemical formula shows the simplest ratio of these ions, and the total charge must always be zero. Example: MgCl₂ (Magnesium Chloride):  ​ Magnesium (Mg²⁺): Loses 2 electrons → +2 charge. ​ Chlorine (Cl⁻): Gains 1 electron → -1 charge. ​ How to balance: You need 2 Cl⁻ ions to cancel Mg²⁺’s +2 charge:+2+(−1×2)=0 (neutral).+2+(−1×2)=0 (neutral). Monatomic Ions ​ A monatomic ion is a single atom with a charge. ○​ Example: Mg²⁺ (magnesium ion), Br⁻ (bromide ion). Oxidation Numbers ​ Oxidation numbers show how many electrons an atom gains or loses in a compound. ○​ Example: Sodium (Na) loses 1 electron → Oxidation number = +1. ○​ Chlorine (Cl) gains 1 electron → Oxidation number = -1. In a compound, the number of electrons lost = number of electrons gained.​ Example: NaCl (Sodium Chloride): ​ Na gives 1 electron (+1), Cl takes it (-1), so the compound is neutral: +1 + -1 = 0. Writing Ionic Formulas 1.​ Cation first, then anion. 2.​ Use subscripts to balance charges. Example: NaF (Sodium Fluoride): ​ Na⁺ (+1) and F⁻ (-1) balance to 0, so the formula is NaF. Polyatomic Ions ​ A polyatomic ion is a group of atoms with a charge (e.g., NH₄⁺, SO₄²⁻). ​ When writing formulas with polyatomic ions, use parentheses if you need more than one. Example: (NH₄)₂O (Ammonium Oxide):  ​ Ammonium (NH₄⁺) has a +1 charge, and oxide (O²⁻) has a -2 charge. ​ You need 2 NH₄⁺ ions to balance 1 O²⁻ ion. Oxyanions ​ Oxyanions are ions made of a nonmetal + oxygen (e.g., NO₃⁻, SO₄²⁻). ​ Nonmetals like nitrogen or sulfur can form different oxyanions. Naming them follows specific rules. Naming Ionic Compounds 1.​ Cations: Use the element name (e.g., Na⁺ = Sodium). 2.​ Anions: Add -ide to the element name (e.g., Cl⁻ = Chloride). 3.​ Cations with multiple charges: Use Roman numerals to show the charge (e.g., Fe²⁺ = Iron(II), Fe³⁺ = Iron(III)). ○​ Example: Fe₂O₃ = Iron(III) Oxide (Fe³⁺, O²⁻). 4.​ Polyatomic Ions: Don’t change their name (e.g., OH⁻ = Hydroxide, SO₄²⁻ = Sulfate). How to Write Formulas from Names 1.​ Identify the cation and anion. 2.​ Balance their charges. 3.​ Use subscripts to show the ratio. Example: Fe₂O₃ (Iron(III) Oxide): ​ Fe³⁺ (Iron with +3 charge), O²⁻ (Oxygen with -2 charge). ​ To balance: 2 Fe³⁺ ions (+6 total) and 3 O²⁻ ions (-6 total). ​ Formula: Fe₂O₃. Lesson 4 Metals aren’t like other materials, but they act like ionic compounds in some ways because they rely on opposites attracting—positive charges and negative charges. Metals in solid form pack together in a tight crowd (a lattice), and each metal atom is surrounded by 8–12 other atoms. Electron sea model=In metals, the atoms don’t really "own" their outermost electrons (the valence electrons). Instead, all these electrons form a “sea” around the metal atoms, moving freely. These freely moving electrons are called delocalized electrons. Since the electrons can move wherever they want, they keep the metal atoms (which are now positive, like little cations) stuck together like glue. This is called a metallic bond. Properties of Metals: Melting and Boiling Points: Metals usually have high melting and boiling points because the glue holding them together (the metallic bond) is strong. But some metals, like mercury, are lazy and melt super easily (that’s why it’s liquid at room temperature). Conductivity:Because those delocalized electrons can move around freely, they’re great at carrying heat and electricity. They’re also why metals are shiny (their electrons like playing with light). Malleability and Ductility: Metals can be squished into sheets (like aluminum foil) or stretched into wires (like copper wires) because their atoms can slide past each other without breaking the glue. The more delocalized electrons a metal has, the tougher it gets. Transition metals like iron are strong because they have extra electrons to share, but soft metals like sodium aren’t as strong. An alloy is a mixture pf elements that has metallic properties. Alloys are stronger, shinier, or more useful than just one metal alone. There are two main types: Substitutional Alloys: This is when some of the original metal atoms are replaced with other metal atoms about the same size. Example: Sterling silver (silver mixed with copper). It’s still shiny but a bit stronger than pure silver. Interstitial Alloys:This is when smaller atoms sneak into the gaps between the bigger metal atoms. Example: Carbon steel (iron with little carbon atoms in the gaps). Carbon makes the iron stronger and harder. ​ Alloys are super useful in everyday life: ​ Stainless steel: Used for sinks and kitchen stuff. ​ Brass: Used for things like candlesticks and musical instruments. ​ Gold alloys: Used in jewelry to make it stronger and less bendy.  Module 7 Lesson 1 Atoms bond to achieve a more stable electronic configuration. Noble gases (Group 18 elements) are naturally stable because their outermost energy levels are filled, making them inert and unlikely to form bonds. Other atoms, however, have incomplete outer shells, which makes them less stable. These atoms can either: Lose, gain, or share electrons to fill or empty their outer energy levels, thereby lowering their potential energy. Reach a state of stability, where their overall energy is minimized. Atoms can form two primary types of bonds to achieve stability: ​ Ionic Bonds: Transfer of electrons. ​ Covalent Bonds: Sharing of electrons. A molecule is formed when two or more atoms bodn covalently. Gaining Stability: The Role of Energy ​ Lower Energy = Higher Stability: Atoms, ions, or molecules with lower potential energy are more stable. Chemical bonding lowers the energy of the system by redistributing valence electrons. ​ Ionic Bonds: Metals lose electrons to become positively charged ions, while nonmetals gain electrons to become negatively charged ions. These oppositely charged ions attract, forming an ionic bond. ​ Covalent Bonds: Atoms share electrons to fill their outermost energy levels. This sharing of electrons allows both atoms to mimic the stable electron configuration of noble gases. A covalent bond forms when two nonmetal atoms share valence electrons. These shared electrons are part of the outer energy levels of both atoms. Some elements naturally exist as molecules of two atoms bonded covalently, like H₂, O₂, and N₂, because the shared electrons make the molecule more stable than the individual atoms. Optimal Bonding Distance= Atoms don't stick too close together because their like charges repel (electrons repel electrons, protons repel protons).​ They also don't stay too far apart because their opposite charges attract (protons of one atom attract the electrons of the other). The perfect distance (optimal bonding distance) happens when Attraction is strongest and Repulsion isn’t too strong. ​ Definition: A single covalent bond involves one pair of shared electrons. ​ Example: Hydrogen Molecule (H₂): Each hydrogen atom has one electron. By sharing one electron with each other, they both achieve the stable configuration of helium (1s²).  ​ Lewis Structure Representation: Single bonds are represented by a line (e.g., H–H) or as a pair of dots (H:H). Group 17 (Halogens)= Halogens like fluorine, chlorine, and bromine form single covalent bonds because they need only one additional electron to complete their octet. Example: Fluorine Molecule (F₂): Each fluorine shares one electron, creating a covalent bond. Group 16 (Oxygen Family)= Elements in this group, like oxygen, can form two covalent bonds because they need two electrons for an octet. Example: Water (H₂O): Oxygen forms two bonds with two hydrogen atoms, completing its octet while each hydrogen achieves helium's configuration. Group 15 (Nitrogen Family)= Elements in this group form three covalent bonds to gain stability. Example: Ammonia (NH₃): Nitrogen forms three covalent bonds with three hydrogen atoms, leaving one lone pair of electrons. Group 14 (Carbon Family)=Carbon forms four covalent bonds because it needs four electrons to fill its outer shell.Example: Methane (CH₄): Carbon shares one electron with each of four hydrogen atoms. Sigma Bonds (σ)forms when the pair of shared electrons is in an area centered between two atoms.It forms when two atomic orbitals overlap directly in a straight line between two atoms. ​ Single bonds (like in H-H or C-H) are sigma bonds. ​ They allow atoms to rotate freely around the bond. ​ In double or triple bonds, only one is a sigma bond, and the rest are weaker bonds called pi bonds. Multiple Covalent Bonds: 1. Double Bonds: Atoms share two pairs of electrons. Triple Bonds: Atoms share three pairs of electrons. 2. Pi Bonds (π): After the first sigma bond, pi bonds form in double and triple bonds by overlapping parallel orbitals. hese bonds exist above and below the axis between the atoms. Orbitals overlap directly in sigma bonds. Parallel orbitals overlap in pi bonds. A single covalent bond is a sigma bond but multiple covalent bonds are made of both sigma and pi bonds. Bond Energy and Reactions:  1.​ Breaking Bonds: Takes energy (endothermic). 2.​ Forming Bonds: Releases energy (exothermic). ​ Endothermic Reactions: Breaking bonds needs more energy than forming new ones. ○​ Example: Breaking water into hydrogen and oxygen (H₂O → H₂ + O₂). ​ Exothermic Reactions: Forming new bonds releases more energy than breaking old ones. ○​ Example: Combustion of methane (CH₄ + O₂ → CO₂ + H₂O). Bond Length:distance between nuclei of bonded atoms. ​ Shorter bonds = stronger bonds. ○​ Single bonds are longest (weakest). ○​ Double bonds are shorter. ○​ Triple bonds are shortest (strongest). ○​ Example: C–C (1.54 Å), C=C (1.34 Å), C≡C (1.20 Å). ​ Larger atoms have longer bonds due to their bigger size. ​ Endothermic reactions : ○​ More energy is used to break bonds than is released when forming new bonds. ○​ Example: H₂O → H₂ + O₂ (energy absorbed). ​ Exothermic reactions : ○​ More energy is released forming bonds than breaking them. ○​ Example: CH₄ + O₂ → CO₂ + H₂O (energy released as heat/light). ○​ Lesson 2 Molecular Compounds and Covalent Bonds: ​ Fluorine (F₂) is a diatomic molecule because each fluorine shares one pair of electrons to achieve a stable electron configuration. ​ Halogens (Group 17, like fluorine and chlorine) have 7 valence electrons and need 1 more electron to complete their octet. They form single covalent bonds with other atoms or with each other (e.g., F₂, Cl₂). ​ Halogens are very reactive in their unbonded state. Naming Covalent Compounds: ​ Prefixes indicate the number of atoms in a compound: ○​ Mono- (1), Di- (2), Tri- (3), Tetra- (4), Penta- (5), etc. ○​ Example: CO₂ = Carbon dioxide (1 carbon, 2 oxygen). ​ Rules:  ○​ Do not use "mono-" for the first element (e.g., CO = Carbon monoxide, not Monocarbon monoxide). ○​ Drop one vowel if two vowels are next to each other (e.g., Monooxide becomes Monoxide). Common Molecular Compound Names: ​ H₂O: Water (Dihydrogen monoxide). ​ NH₃: Ammonia. ​ N₂O: Nitrous oxide (laughing gas). ​ NaCl: Table salt (Sodium chloride). ​ NaHCO₃: Baking soda (Sodium hydrogen carbonate). Acids and Their Names: 1.​ Binary Acids (Hydrogen + 1 other element): ○​ Name: "Hydro-" + root + "-ic" + "acid." ○​ Example: HCl = Hydrochloric acid. 2.​ Oxyacids (Hydrogen + Polyatomic Ion with Oxygen): ○​ Name based on the polyatomic ion. ○​ Example: HNO₃ (contains nitrate) = Nitric acid. Exothermic and Endothermic Reactions: ​ Exothermic: Release energy when bonds form (e.g., combustion). ​ Endothermic: Absorb energy to break bonds (e.g., photosynthesis). Writing Formulas from Names: 1.​ Use prefixes to determine the number of each atom. ○​ Example: Dinitrogen tetroxide = N₂O₄. 2.​ For acids: ○​ Binary acids: Start with hydrogen + 1 element. ​ Example: Hydrochloric acid = HCl. ○​ Oxyacids: Use the polyatomic ion. ​ Example: Nitric acid (nitrate ion) = HNO₃.  Lesson 3 Structural Formulas: ​ Definition: Show how atoms are arranged and bonded in a molecule using symbols and lines (bonds). ​ Purpose: Provide more detail than molecular formulas by showing the bonding arrangement. ​ Models: ○​ Ball-and-stick: Atoms = spheres, bonds = sticks. ○​ Space-filling: Shows atom size and arrangement without bonds. ○​ Lewis Structures: 2D diagrams showing bonds and electron pairs. How to Draw Lewis Structures: 1.​ Count Valence Electrons: Add up electrons for all atoms. 2.​ Central Atom: Place the least electronegative atom in the center (except hydrogen). 3.​ Distribute Electrons: Give atoms 8 electrons (octet rule); hydrogen gets only 2. 4.​ Use Multiple Bonds: If needed, use double or triple bonds to satisfy octets. ○​ For Ions: ​ Add electrons for negative charges. ​ Subtract electrons for positive charges. Resonance Structures: ​ Definition: Some molecules/ions can be drawn in multiple valid ways, differing only in the position of electron pairs. ​ Example: Nitrate ion (NO₃⁻) has 3 resonance structures with one double bond and two single bonds rotating among oxygens. ​ Key Point: The actual structure is an average of the resonance forms, leading to identical bond lengths (intermediate between single and double bonds). Exceptions to the Octet Rule: 1.​ Odd Valence Electrons: ○​ Example: Nitrogen dioxide (NO₂) has 17 electrons, leaving nitrogen with only 7 in its outer shell. 2.​ Suboctets (Less Than 8 Electrons): ○​ Example: Boron trifluoride (BF₃) has only 6 electrons around boron, making it very reactive. 3.​ Coordinate Covalent Bonds: ○​ One atom donates both electrons for a bond. ○​ Example: In ammonium (NH₄⁺), nitrogen donates a lone pair to bond with H⁺. 4.​ Expanded Octets (More Than 8 Electrons): ○​ Possible for atoms in Period 3 or higher (e.g., phosphorus, sulfur) because they have d-orbitals. ○​ Examples: Phosphorus pentachloride (PCl₅) and sulfur hexafluoride (SF₆). Here are your questions in a randomized order: 1. The periodic law states that the physical and chemical properties of the elements are periodic function of their atomic a. masses b. numbers c. radii d. charges 2. Which of the following is an intensive property of matter? a. amount of energy b. density c. volume d. mass 3. An atom is a. the smallest unit of matter that maintains its chemical identity b. the smallest unit of a compound c. always made of carbon d. smaller than an electron 4. The equation E = hv helped Louis de Broglie determine a. how protons and neutrons behave in the nucleus b. how electron wave frequencies correspond to specific energies  c. whether electrons behave as particles d. whether electrons exist in a limited number of orbits with different energies 5. To which block do the actinide elements belong? a. d block b. s block c. f block d. p block 6. A neutral group of atoms held together by covalent bonds is a a. molecular formula b. chemical formula c. polyatomic ion d. molecule 7. The principle that states that the physical and chemical properties of the elements are periodic function of their atomic numbers is a. the periodic table b. the periodic law c. the law of properties d. Mendeleev’s law 8. The electrons involved in the formation of a chemical bond are called a. dipoles b. s electrons  c. Lewis electrons d. valence electrons 9. Which of the following is not part of Dalton’s atomic theory? a. atoms cannot be divided, created, or destroyed b. the number of protons in an atom is its atomic number c. in chemical reactions, atoms are combined, separated, or rearranged d. all matter is composed of extremely small particles called atoms 10. A horizontal row of blocks in the periodic table is called a(n) a. group b. period c. family d. octet 11. Dalton incorporated the law of conservation of mass into his atomic theory by asserting that a. atoms are indivisible b. atoms of different elements have different properties c. matter is composed of atoms d. atoms can be destroyed in chemical reactions 12. Bond length is the average distance between two bonded atoms a. at which potential energy is at a minimum b. at which kinetic energy is at a maximum  c. at which potential energy is at a maximum d. and equal to one-half the diameter of the electron cloud 13. The law of definite proportions a. contradicted Dalton’s atomic theory b. agrees with Dalton’s atomic theory c. replaced the law of conservation of mass d. assumes that atoms of all elements are identical 14. The law of conservation of mass a. contradicted Dalton’s atomic theory b. agrees with Dalton’s atomic theory c. replaced the law of conservation of energy d. assumes that atoms of all elements are identical 15. The set of orbitals that are dumbbell-shaped and directed along the x, y, and z axes are called a. d orbitals b. p orbitals c. f orbitals d. s orbitals 16. A chemical can be defined as a. a toxic substance b. an unnatural additive placed in food  c. any substance that has a definite composition d. any substance that is not alive 17. Which of the following is an extensive property of matter? a. melting point b. boiling point c. volume d. density 18. The idea of arranging the elements in the periodic table according to their chemical and physical properties is attributed to a. Mendeleev b. Moseley c. Bohr d. Ramsay 19. When atoms share electrons, the electrical attraction of an atom for the shared electrons is called the atom’s a. electron affinity b. electronegativity c. resonance d. hybridization 20. Mendeleev predicted that the spaces in his periodic table represented a. isotopes  b. radioactive elements c. unstable elements d. undiscovered elements 21. If two covalently bonded atoms are identical, the bond is a. nonpolar covalent b. polar covalent c. dipole covalent d. coordinate covalent 22. A mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together is called a(n) a. dipole b. chemical bond c. Lewis structure d. London force 23. Nonpolar covalent bonds are not common because a. one atom usually attracts electrons more strongly than the other b. ions always form when atoms join c. the electrons usually remain equally distant from both atoms d. dipoles are rare in nature 24. The electron configuration of aluminum, atomic number 13, is [Ne] 3s² 3p¹. Aluminum is in Period  a. 2 b. 3 c. 6 d. 13 25. The principle that states that the physical and chemical properties of the elements are periodic function of their atomic numbers is a. the periodic table b. the periodic law c. the law of properties d. Mendeleev’s law 26. A state of matter in which a material has no definite shape but has a definite volume is the ____ state a. gas b. liquid c. plasma d. solid 27. The electrostatic attraction between positively charged nuclei and negatively charged electrons permits two atoms to be held together by a(n) a. chemical bond b. London force c. neutron d. ion 28. The two most important properties of all matter are a. the ability to carry an electric current well and to hold electric charge b. taking up space and having mass c. being brittle and hard d. being malleable and ductile 29. Calcium, atomic number 20, has the electron configuration [Ar] 4s². In what period is calcium? a. Period 2 b. Period 4 c. Period 8 d. Period 20 30. The greater the electronegativity difference between two bonded atoms, the greater the percentage of ____ character a. ionic b. covalent c. metallic d. electron sharing 31. For an electron in an atom to change from the ground state to an excited state, a. energy must be released b. energy must be absorbed c. radiation must be emitted d. the electron must make a transition from a higher to a lower energy level 32. If the atoms that share electrons have an unequal attraction for the electrons, the bond is called a. nonpolar b. polar c. ionic d. dipolar 33. If 6.0 g of element K combine with 17 g of element L, how many grams of element K combine with 85 g of element L? a. 17 g b. 23 g c. 30 g d. 91 g 34. A quantum of electromagnetic energy is called a(n) a. photon b. electron c. excited atom d. orbital 35. The law of definite proportions a. contradicted Dalton’s atomic theory b. agrees with Dalton’s atomic theory c. replaced the law of conservation of mass  d. assumes that atoms of all elements are identical 36. The B—F bond in BF₃ (electronegativity for B is 2.0; electronegativity for F is 4.0) is a. polar covalent b. ionic c. nonpolar covalent d. metallic 37. The periodic law that allows some properties of an element to be predicted based on its a. position in the periodic table b. number of isotopes c. symbol d. color 38. According to Dalton’s atomic theory, atoms a. are destroyed in chemical reactions b. can be divided c. of each element are identical in size, mass, and other properties d. of different elements cannot combine 39. The atomic number of an element is the number of a. protons b. electrons c. neutrons  d. isotopes 40. In oxides of nitrogen, such as N₂O, NO, NO₂, and N₂O₃, atoms combine in small whole-number ratios. This evidence supports the law of a. conservation of mass b. definite composition c. multiple proportions d. mass action 41. If each atom of element D has 3 mass units and each atom of element E has 5 mass units, a molecule composed of one atom each of D and E has a. 2 mass units b. 8 mass units c. 15 mass units d. 35 mass units 42. The set set of orbitals that are dumbbell-shaped and directed along the x, y, and z axes are called a. d orbitals b. p orbitals c. f orbitals d. s orbitals 43. Which is not part of hydrogen’s line-emission spectrum? a. Balmer series b. Lyman series  c. Paschen series d. Aufbau series 44. The law of conservation of mass states that the mass of the reactants in a chemical reaction equals the mass of the products. a. True b. False 45. The electrons in the outermost shell of an atom are called a. inner electrons b. core electrons c. valence electrons d. excited electrons 46. Mendeleev left spaces in his periodic table and predicted the existence of three elements and their a. atomic numbers b. colors c. properties d. radioactivity 47. The period of an element can be determined from its a. reactivity b. density c. symbol  d. electron configuration 48. A chemical bond is formed by the attraction between a. an ion and a molecule b. an electron and a nucleus c. a proton and a neutron d. the nuclei and electrons of two atoms 49. A state of matter in which a material has no definite shape but has a definite volume is the ____ state a. gas b. liquid c. plasma d. solid 50. If 4.0 g of element A combine with 10 g of element B, then 12 g of element A combine with ____ g of element B. a. 10 b. 12 c. 24 d. 30 51. According to the law of conservation of mass, when sodium, hydrogen, and oxygen react to form a compound, the mass of the compound is ____ the sum of the masses of the individual elements a. equal to  b. greater than c. less than d. either greater than or less than 52. The electron in a hydrogen atom has its lowest total energy when the electron is in its a. neutral state b. excited state c. ground state d. quantum state 53. Basic research is usually performed a. to develop new products b. to understand an environmental problem c. to gain knowledge d. to solve a particular problem 54. If 4.0 g of element A combine with 10 g of element B, then 12 g of element A combine with ____ g of element B. a. 10 b. 12 c. 24 d. 30 55. The idea of arranging the elements in the periodic table according to their chemical and physical properties is attributed to a. Mendeleev b. Moseley c. Bohr d. Ramsay