Inorganic Chemistry L1: Matter PDF

(In)organic Chemistry L1: Matter “All everyday objects that can be touched are ultimately composed of atoms.” Matter Matter is anything that occupies space and has mass/volume Matter is made up of atoms and particles. Matter can be synthesized or occur naturally Matter can...

(In)organic Chemistry L1: Matter “All everyday objects that can be touched are ultimately composed of atoms.” Matter Matter is anything that occupies space and has mass/volume Matter is made up of atoms and particles. Matter can be synthesized or occur naturally Matter can be broken down into two categories: ○ Pure Substances — A single kind of matter that cannot be separated into other kinds of matter; same through, made by the same material. Elements → Consist of only one atom (carbon, gold, silver) Compound → Chemically combining two or more elements. (H20, NaCl) ○ Mixtures — Composed of more than one component that is separable (not chemically bonded). Homogeneous has a uniform composition and properties throughout. The components are evenly distributed and cannot be distinguished from one another. Heterogeneous Has a non-uniform composition and properties. The components are not evenly distributed and can be distinguished from one another. Inorganic Compound (acids, bases, salts) Dissociate in solution (water) which forms charged particles/ion such as:  Cation → positively charged ion Anion → negatively charged ion Acid ○ All acids contain hydrogen atoms in their chemical structure, which can be released as hydrogen ions (H+) when the acid is dissolved in water. Sulfuric acid, citric acid, nitric acid. Base ○ Contains the compound: hydroxide (OH) in their chemical structure which releases (OH-) when the base is dissolved in water. antacid tablets, laundry detergents, soaps and other bath products, baking soda, bleach Salt ○ Neutralization; the reaction between acid and base which produces table salt (NaCl). potassium dichromate, calcium chloride, sodium bisulfate, and copper sulfate. Covalent Bond Water (H20) Electrons are not equally shared within the molecule as they have a higher probability as they have the higher probability of them being located in the oxygen atom. Giving one end of the molecule a slightly negative charge (yung right side ng oxygen atom) and the other end (kung nasaan yung H dun yung positive charge.)  Hydrogen bond yung na aattract yung postive pole kay negative pole. Types of matter can be distinguished through 2 components: Composition → refers to the different components of matter along with their relative proportion. Properties → refers to the qualities/attributes that distinguish one sample of matter from another. Has 2 categories: ○ Physical Properties – Without change in chemical composition–density, mass, shape, height, length, freezing, condensation, evaporation, Cutting, tearing, shattering, hammering, mixing. ○ Temporary changes Intensive properties Extensive properties ○ Chemical Properties Permanent changes; transform substances into entirely new substances–oxidation, combustion Involve changes in chemical composition and properties. Matter exist in three fundamental states namely: Solid: Definite volume and shape. Maintains shape regardless of the container. Particles are closely packed in a regular three-dimensional array. Strong forces between particles prevent free movement. Can change shape only by force (e.g., smashing or squashing). Liquid: Definite volume but indefinite shape.  Takes the shape of the container it occupies. Particles are close together but can move around and slide past one another. Slightly compressible. Gas: No definite shape or volume. Particles move randomly and are widely separated. Expands to fill the entire volume of the container. Highly compressible due to large gaps between particles. Phase Change https://chem.libretexts.org/Bookshelves/Inorganic_Chemistry/Supplemental_Modules _and_Websites_(Inorganic_Chemistry)/Chemical_Reactions/Properties_of_Matter The transformation of matter or the chemical reactions are governed by certain basic rules or laws, namely: the law of conservation of mass, the law of definite proportions, and the law of multiple proportions. The Law of Conservation of Mass states that the total mass of the reactants in a chemical reaction is equal to the total mass of the products. In other words, matter is neither created nor destroyed during a chemical reaction, it is only transformed. Proposed by French chemist Antoine-Laurent Lavoisier in the 18th century States that the total mass before a reaction = total mass after a reaction Explains that atoms are rearranged but not created or destroyed in a chemical reaction “Mass is neither created nor destroyed during a chemical transformation.” The Law of Definite Proportion, also known as the Law of Constant Composition, states that a chemical compound always contains the same elements combined in the same fixed proportion by mass, regardless of the source or method of preparation. Proposed by French chemist Joseph Proust in the late 18th century States that the composition of a pure chemical compound is always the same The relative amounts of the elements in a compound are fixed and do not vary ○ For example, water (H2O) will always be composed of 2 hydrogen atoms and 1 oxygen atom, regardless of the source of the water. The relative amounts of hydrogen and oxygen in water do not change. The Law of Multiple Proportion states that when two elements form more than one compound, the masses of one element that combine with a fixed mass of the other element are in a ratio of small whole numbers. Proposed by English chemist John Dalton in the early 19th century Explains that elements can combine in different fixed ratios to form different compounds The ratios of the masses of the elements in these compounds are simple whole number ratios For example, carbon and oxygen can form two different compounds: Carbon monoxide (CO) has a mass ratio of 12:16 (carbon to oxygen) Carbon dioxide (CO2) has a mass ratio of 12:32 (carbon to oxygen) These ratios are simple whole number ratios of 1:1 and 1:2 respectively. (In)organic Chemistry L2: Atomic Theory Atomic Theory Originated from Greek and India Describes the nature of atoms States that matter is made of tiny particles called atoms Atoms are made up of subatomic particles Atom Democritus proposed that matter is made up of tiny indestructible and indivisible units called atoms Came from the greek word: Atomos, which means indivisible Its center is called the nucleus. Atomism Suggests that matter is composed of discrete particles Early explanation for the nature of matter and is not based on empirical data John Dalton’s Atomic Model Proposed that all matter is composed of atoms This model suggested that atoms are circular in shape J.J. Thomson’s Atomic model (1897) Discovered electrons Realized that atoms are made up for even smaller particles Atomic nucleus is not yet discovered. Plum-pudding method proposed in 1904, exhibits that atoms are made up of negative electrons that float in a “soup” of positive charge, much like plums in a pudding or raisins in a fruit cake. (parang sopas, pasta si negative electron, tas yung positive charges is yung sabaw.) Radiation (Marie and Pierre Curie [20th century]) Discovered that radioactive elements emit particles which are capable of passing through matter, similarly with X-rays. Ernest Rutherford’s Atomic Model His model described the atoms as tiny, dense, positively charged core (nucleus) surrounded by lighter, negatively charged electrons Could not explain why atoms only emit light at certain wavelengths or frequencies Nuclear model Gold foil experiment In 1920, he predicted that another kind of particle is present along the nucleus with the proton Proton Niels Bohr Solved Ernest Rutherford’s oversight by proposing that electrons could only orbit the nucleus on certain special orbits at different energy levels around the nucleus. Planetary Model Electron shell James Chadwick In 1920, Rutherford predicted the existence of a new particle in the nucleus, in addition to the protons. Rutherford reasoned that if only positively charged protons were present in the nucleus, it would break apart due to repulsive forces. To maintain electrical neutrality in the atom, this additional particle must be neutral. In 1932, James Chadwick discovered the neutron and its mass. Lousie the bukol ○ Fundamental Particles of Atoms Subatomic particles are smaller than the atom. The three main subatomic particles that form the atom are: ○ Protons (Ernest Rutherford, 1919)  Conducted the gold foil experiment; where he projected alpha particles (helium nuclei) at a gold foil, and the positive alpha particles were deflected. Concluded that protons exist in a nucleus and have a positive nuclear charge. Number of protons is the atomic number of an element. Are used to determine an element ○ Electron (e-) Are located in an electron electron cloud; an area surrounding the nucleus Electrons are closer to the nucleus of the atom Are negatively charged subatomic particles equal to the magnitude of the positive charges of protons Has no mass–meaning it is lighter than proton and neutron Were discovered by: John Joseph / J.J. Thomson, 1897 ○ Demonstrated the ratio of mass to the electric charge of the cathode rays (electrons) ○ Confirmed that cathode rays (electrons) are fundamental particles that are negatively charged Robert Milikan ○ Found the value of the electronic charge via oil drop experiment. Unequal amount of proton and electron create ions may form: Protons > Electrons → Cation (positive ion) Electrons > Protons → Anion (negative ion) ○ > = greater than (WALA LANG BWAHAHAHAHA) ○ Neutron (James Chadwick, 1932) Are located in the nucleus with the protons making up the mass of the atom Demonstrated penetrating radiation consisted of beams of neutral particles Neutron number can be found by subtracting proton from the number of atomic mass Determines the isotope of an atom and its stability. Number of neutrons is not equal to the number of protons Properties of Atoms Solid, liquid, gas, and plasma is composed of neutral or ionized atoms The chemical properties of an atom are determined by: ○ the number of protons ○ number of electrons and its arrangement–following the principles of quantum mechanics. The number of electron in the outermost valence shell (electron shell) determines the chemical bonding behavior Atomic Number or Proton number (Z) The total number of protons in the nucleus For an element to be electrically neutral, it must have equal amount of protons and electrons The nuclear charge is +Ze, where Z is the atomic number and e is the elementary charge. Elementary charge (e) ○ The value of the electric charge carried by a single proton or electron. ○ It is a constant with a value of 1.602 x 10-19 coulombs (C) Nuclear charge or Total amount of proton or the value of the atomic number ○ Atomic number (Z) is the number of protons in the nucleus of an atom; which is equal to the nuclear charge, which is the total positive charge of the nucleus. ○ The nuclear charge is what attracts the electrons to the atom. Within an atom, the positive nuclear charge creates an electric field that extends outward from the nucleus. This electric field influences the motion and behavior of the negatively charged electrons orbiting the nucleus. The  strength of the electric field decreases with increasing distance from the nucleus. Each electron is influenced by the electric fields produced by the positive nuclear charge and the other (Z – 1) negative electrons in the atom. Pauli Exclusion Principle Requires the electrons in an atom to occupy different energy levels instead of condensing in the ground state (the lowest possible energy level–where the electrons occupy the orbitals closest to the nucleus) The ordering of the electrons in the ground state of multielectron atoms, moves progressively up the energy scale until each of the atom’s electrons has been assigned a unique set of quantum numbers Mass Number - Mass of atom (A) Since electron are almost massless in comparison to nucleons (a single proton or neutron), the total number of protons and neutrons in the nucleus of an atom determines the atomic mass. The mass number is different for each different isotopes of a chemical element. Mass number is written either after the element name or as a superscript to the left of an element’s symbol. ○ The most commom isotope of carbon is carbon-12 or 12C The unit of measure for the mass of an atom is atomic mass unit (amu or u) and it has a constant of 1.66 x 10^-24 Amu is the second mass standard, next to Kg One unified atomic mass unit is approximately the mass of 1 nucleon (either a single proton or neutron / nucleus) and is numerically equivalent to 1 g/mol For 12C the atomic mass is exactly 12u, since the atomic mass unit is defined from it. For other isotopes, the isotopic mass usually differs and is usually within 0.1 u of the mass number. For example, 63Cu (29 protons and 34 neutrons) has a mass number of 63 and an isotopic mass in its nuclear ground state is 62.91367 u. There are two reasons for the difference between mass number and isotopic mass, known as the mass defect: 1. The neutron is slightly heavier than the proton. This increases the mass of nuclei with more neutrons than protons relative to the atomic mass unit scale based on 12C with equal numbers of protons and neutrons. 2. The nuclear binding energy 3. Links to an external site. 4. varies between nuclei. A nucleus with greater binding energy has a lower total energy, and therefore a lower mass according to Einstein’s mass-energy equivalence 5. relation E = mc2. For 63Cu the atomic mass is less than 63 so this must be the dominant factor. Electron configuration  - describes how electrons are distributed in its atomic orbitals. It follow a standard notation in which all electron-containing atomic subshells (with the number of electrons they hold written in superscript) are placed in a sequence. For example, the electron configuration of sodium is 1s22s22p63s1. However, the standard notation often yields lengthy electron configurations (especially for elements having a relatively large atomic number). In such cases, an abbreviated or condensed notation may be used instead of the standard notation. In the abbreviated notation, the sequence of completely filled subshells that correspond to the electronic configuration of a noble gas is replaced with the symbol of that noble gas Links to an external site. in square brackets. Therefore, the abbreviated electron configuration of sodium is [Ne]3s1 (the electron configuration of neon is 1s22s22p6, which can be abbreviated to [He]2s22p6). Electron Configurations are useful for: Determining the valency of an element.  Predicting the properties of a group of elements (elements with similar electron configurations tend to exhibit similar properties). Interpreting atomic spectra. This notation for the distribution of electrons in the atomic orbitals of atoms came into practice shortly after the Bohr model of the atom was presented by Ernest Rutherford and Niels Bohr in the year 1913. Writing Electron Configurations Shells The maximum number of electrons that can be accommodated in a shell is based on the principal quantum number (n). It is represented by the formula 2n2, where ‘n’ is the shell number. The shells, values of n, and the total number of electrons that can be accommodated are tabulated below. Subshells The subshells into which electrons are distributed are based on the azimuthal quantum number Links to an external site. (denoted by ‘l’). This quantum number is dependent on the value of the principal quantum number, n. Therefore, when n has a value of 4, four different subshells are possible. When n=4. the subshells correspond to l=0, l=1, l=2, and l=3 and are named the s, p, d, and f subshells respectively. The maximum number of electrons that can be accommodated by a subshell is given by the formula 2*(2l + 1). Therefore, the s, p, d, and f subshells can accommodate a maximum of 2, 6, 10, and 14 electrons respectively. All the possible subshells for values of n up to 3 are tabulated below Thus, it can be understood that the 1p, 2d, and 3f orbitals do not exist because the value of the azimuthal quantum number is always less than that of the principal quantum number. Notation The electron configuration of an atom is written with the help of subshell labels. These labels contain the shell number (given by the principal quantum number), the subshell name (given by the azimuthal quantum number), and the total number of electrons in the subshell in superscript. For example, if two electrons are filled in the ‘s’ subshell of the first shell, the resulting notation is ‘1s2’. With the help of these subshell labels, the electron configuration of magnesium (atomic number 12) can be written as 1s2 2s2 2p6 3s2. Filling of Atomic Orbitals Aufbau Principle This principle is named after the German word ‘Aufbeen’ which means ‘build up’.  The Aufbau principle Links to an external site. dictates that electrons will occupy the orbitals having lower energies before occupying higher energy orbitals. The energy of an orbital is calculated by the sum of the principal and the azimuthal quantum numbers. According to this principle, electrons are filled in the following order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p… The order in which electrons are filled in atomic orbitals as per the Aufbau principle is illustrated below. It is important to note that there exist many exceptions to the Aufbau principle such as chromium and copper. These exceptions can sometimes be explained by the stability provided by half-filled or completely filled subshells. Pauli's Exclusion Principle The Pauli exclusion principle states that a maximum of two electrons, each having opposite spins, can fit in an orbital.  This principle can also be stated as “no two electrons in the same atom have the same values for all four quantum numbers”. Therefore, if the principal, azimuthal, and magnetic numbers are the same for two electrons, they must have opposite spins. Hund’s Rule This rule describes the order in which electrons are filled in all the orbitals belonging to a subshell. It states that every orbital in a given subshell are singly occupied by electrons before a second electron is filled in an orbital. In order to maximize the total spin, the electrons in the orbitals that only contain one electron all have the same spin (or the same values of the spin quantum number). An illustration detailing the manner in which electrons are filled in compliance with Hund’s rule of maximum multiplicity is provided above. Examples The electron configurations of a few elements are provided with illustrations in this subsection. Electron Configuration of Hydrogen The atomic number Links to an external site. of hydrogen is 1. Therefore, a hydrogen atom contains 1 electron, which will be placed in the s subshell of the first shell/orbit. The electron configuration of hydrogen is 1s1, as illustrated below. Electron Configuration of Hydrogen Electron Configuration of Oxygen The atomic number of oxygen is 8, implying that an oxygen atom holds 8 electrons. Its electrons are filled in the following order: K shell – 2 electrons L shell – 6 electrons Therefore, the electron configuration of oxygen is 1s2 2s2 2p4, as shown in the illustration provided below. Electron Configuration of Oxygen Chlorine Electronic Configuration Chlorine has an atomic number of 17. Therefore, its 17 electrons are distributed in the following manner: K shell – 2 electrons L shell – 8 electrons M shell – 7 electrons The electron configuration of chlorine is illustrated below. It can be written as 1s22s22p63s23p5 or as [Ne]3s23p5 Electron Configuration of Chlorine Please Remember the following: The electronic configuration of an element is a symbolic notation of the manner in which the electrons of its atoms are distributed over different atomic orbitals. While writing electron configurations, a standardized notation is followed in which the energy level and the type of orbital are written first, followed by the number of electrons present in the orbital written in superscript. For example, the electronic configuration of carbon (atomic number: 6) is 1s22s22p2. The three rules that dictate the manner in which electrons are filled in atomic orbitals are:  The Aufbau principle: electrons must completely fill the atomic orbitals of a given energy level before occupying an orbital associated with a higher energy level. Electrons occupy orbitals in the increasing order of orbital energy level. Pauli’s exclusion principle: states that no two electrons can have equal values for all four quantum numbers. Consequently, each subshell of an orbital can accommodate a maximum of 2 electrons and both these electrons MUST have opposite spins. Hund’s rule of maximum multiplicity: All the subshells in an orbital must be singly occupied before any subshell is doubly occupied. Furthermore, the spin of all the electrons in the singly occupied subshells must be the same (in order to maximize the overall spin) Electron configurations provide insight into the chemical behaviour of elements by helping determine the valence electrons of an atom. It also helps classify elements into different blocks (such as the s-block elements, the p-block elements, the d-block elements, and the f-block elements). This makes it easier to collectively study the properties of the elements Quantum numbers are used to describe the location of an electron in an associated atom. It specifies the properties of the atomic orbitals and the electrons in those orbitals. An electron in an atom or ion has four quantum numbers to describe its state. Think of them as important variables in an equation which describes the three-dimensional position of electrons in a given atom. There are four quantum numbers: a. n - principal quantum number Links to an external site. : describes the energy level b. ℓ - azimuthal  Links to an external site. or angular momentum quantum number Links to an external site. : describes the subshell c. mℓ or m - magnetic quantum number: describes the orbital of the subshell d. ms or s - spin quantum number Links to an external site. : describes the spin For further information regarding the quantum numbers, please click the link below: https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_ Textbook_Maps/Supplemental_Modules_(Physical_and_Theoretical_Chemistry )/Quantum_Mechanics/10%3A_Multi-electron_Atoms/Quantum_Numbers_for_ Atoms#:~:text=The%20principal%20quantum%20number%2C%20n,the%20sha pe%20of%20the%20orbital. Links to an external site. Shape of Orbitals s Orbitals For any value of n, a value of l = 0 places that electron in an s orbital. This orbital is spherical in shape: p Orbitals There are three possible orbitals when l = 1. These are designated as p orbitals and have dumbbell shapes. Each of the p orbitals has a different orientation in three-dimensional space. d Orbitals When l = 2, m1 values can be −2, −1, 0, +1, +2 for a total of five d orbitals. Note that all five of the orbitals have specific three-dimensional orientations. f Orbitals The most complex set of orbitals are the f orbitals. When l = 3, ml values can be −3, −2, −1, 0, +1, +2, +3 for a total of seven different orbital shapes. Again, note the specific orientations of the different f orbitals.