General Chemistry I PDF

General Chemistry I anything that occupy space and has mass Volume Mass/ Weight  SOLID LIQUID GAS In a solid, the particles Atoms have many The spaces between gas (ions, atoms or molecules) nearest...

General Chemistry I anything that occupy space and has mass Volume Mass/ Weight  SOLID LIQUID GAS In a solid, the particles Atoms have many The spaces between gas (ions, atoms or molecules) nearest neighbors in molecules are very big. are closely packed contact, yet no long-range together. order is present. PLASMA BOSE-EINSTEIN FERMIONIC In a plasma, electrons are CONDENSATE CONDENSATE ripped away from their typically formed when a It is a superfluid phase nuclei, forming an electron gas of bosons at very low formed by fermionic “sea”. This gives it the densities is cooled to particles at low ability to conduct temperatures very close temperatures. electricity. to absolute zero  MATTER PURE SUBSTANCE MIXTURE ELEMENTS COMPOUNDS HOMOGENEOUS HETEROGENEOUS METALS SOLUTION COLLOID NON-METALS SUSPENSION METALLOIDS A homogeneous mixture of one or more solutes dissolved in a solvent Example: Water and salt  Chromatography is the separation of a mixture by passing it in solution or suspension, or as a vapor (as in gas chromatography), through a medium in which the components move at different rates.  Distillation is a purification process where the components of a liquid mixture are vaporized and then condensed and isolated.  Evaporation is a technique used to separate out homogeneous mixtures that contain one or more dissolved salts. Filtration is a separation method used to separate out pure substances in mixtures comprised of particles—some of which are large enough in size to be captured with a porous material.  a process by which components in a chemical mixture are separated into different parts (called fractions) according to their different boiling points. The basic unit of matter.  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 440 BCE Proposed that everything in the world is made up of tiny particles. He imagined that if you keep on dividing things you would eventually reach to a point that you cannot cut it anymore. He called the particles “atomos” which means “indivisible”.  Atomic Theory 01 5th Century BC Atomism According to atomism, all matter is made up of tiny, indivisible particles called atoms, which are in constant motion and are too small to be seen with the naked eye.  John Dalton 1808 Dalton’s Atomic Theory 1. Each chemical element is composed of extremely small particles that are indivisible and cannot be seen by the naked eye, called atoms. 2. All atoms of an element are alike in mass and other properties, but the atoms of one element differ from all other elements. 3. For each compound, different elements combine in a simple numerical ratio.  Atomic Theory 02 Solid Sphere Model 1808 According to this model, atoms are tiny, indestructible spheres with no internal structure. Solid Sphere Model John Dalton of the Atom If matter were composed of atoms, what were atoms composed of? Were they the smallest particles, or was there something smaller? Joseph John Thompson 1897 By using the cathode ray experiment he discovered the ELECTRONS. He also found out that the particle that make up the cathode ray is 1000 times smaller that a Hydrogen atom.  Atomic Theory 03 Plum Pudding Model electron 1904 J.J. Thompson proposed that the atom is composed of a positively charged sphere with negatively charged electrons distributed throughout it. J.J. Thomson's Plum Pudding Model of the Atom The negatively charged electrons were embedded in a positively charged "pudding" of matter, which made up most of the atom's mass.  Ernest Rutherford 1909 He conducted the Gold Foil Experiment which led to the discovery of PROTONS. He suggested that a neutrally charged particle, also resided in the nuclei of atoms and coined the term “NEUTRONS”.  James Chadwick 1932 He fired alpha radiation at beryllium sheet from a polonium source that produces an uncharged, penetrating radiation. His experiment led to the discovery of “Neutrons  Atomic Theory 04 Nuclear Model 1911 electron Ernest Rutherford proposed that most of the mass of the atom is concentrated in a tiny, positively charged nucleus at the center, with negatively charged nucleus electrons orbiting around it. Ernest Rutherford’s Nuclear This model became the basis for our Model of the Atom understanding of atomic structure today. Niels Bohr 1913 By drawing on the earlier works of Max Planck and Albert Einstein he improved the Nuclear Model. He stipulated that electrons orbit the nucleus at fixed energies and distances.  Atomic Theory 04 Planetary Model 1911 According to the Bohr model, often referred to as a planetary model, the electrons encircle the nucleus of the atom in specific allowable paths called orbits. When the electron is in one of these orbits, its energy is fixed.  Werner Heisenberg 1927 Heisenberg Uncertainty Principle states that we cannot know both the position and speed of a particle, such as a photon or electron, with perfect accuracy. the more we nail down the particle's position, the less we know about its speed and vice versa. Erwin Schrödinger 1927 famous for his Schrodinger’s Cat experiment. formulated the Schrodinger’s equation, a fundamental equation in quantum mechanics that gave rise to the Quantum Mechanical Model of Atoms.  Atomic Theory 05 Quantum Model 1920s The quantum mechanical model uses these wave functions. It helps explain atomic and molecular structures. It predicts electron behavior more accurately. QUESTIONS? THANK YOU! INTRODUCTION TO ELECTRON CONFIGURATION  The main or principal energy levels (n) are numbered, starting with n= 1 as the energy level nearest to the nucleus and going to n=7. The maximum number of electrons for a specific energy level can be calculated from the formula 2n2 Principal Energy Level, n Maximum No. of Electrons Allowed per Energy Level = 2n2 1 2 x (1)2 = 2 2 2 x (2)2 = 8 3 2 x (3)2 = 18 4 2 x (4)2 = 32 EXERCISE: 1. Ge 2. Si  EXERCISE: 1. Ge = e- 32 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p2  EXERCISE: 1. Ge = e- 32 1s 2s 2p 3s 3p 4s 3d 4p 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p2 ELECTRON CONFIGURATION EXERCISE: 2. Si = e- 14 1s2 2s2 2p6 3s2 3p2  EXERCISE: 2. Si = e- 14 1s 2s 2p 3s 3p 1s2 2s2 2p6 3s2 3p2 ELECTRON CONFIGURATION  ELECTRON CONFIGURATION  What is Electron Configuration? It describes how electrons are distributed in its atomic orbitals. It follows 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.  The arrangement of electrons must obey the following three rules: Aufbau Principle This principle is named after the German word ‘Aufbeen’ which means ‘build up’. It 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 arrangement of electrons must obey the following three rules: Pauli 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”.  The arrangement of electrons must obey the following three rules: 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 is singly occupied by electrons before a second electron is filled in an orbital. Drawing Electron Configuration Diagrams * * *Proton: Number of e- and protons are the same *Neutron: Atomic mass minus Atomic number  Example: K: ✓ e- = ✓ p+ = ✓ n = ✓ Electron Configuration:  Example: K: ✓ e- = 19 ✓ p+ = 19 ✓ n =20 ✓ Electron Configuration: 1s2 2s2 2p6 3s2 3p6 4s1  Let’s practice: 1. Li 2. C 3. Si 4. Ge Thank you! Question? Quantum Numbers & Shapes of Orbitals Kristin Shane Sumayao-Morco General Chemistry I Drawbacks of Bohr’s Atomic Model Objections were being made on Bohr’s atomic model about: - Theory of mixture of classical and quantum physics. - Model could not be used to explain variuos intensities and some spectral lines.  He formulated an equation called In 1926,he gave “the Schrödinger the idea that of equation”, in wave motion of 01 03 which electrons electrons are treated as moving with wave like motion in 3D space around the nucleus. He won the The solution of Schrödinger Nobel Prize in 02 05 Equation gave a Physics in 1933 set of numerical Erwin Rudolf Schrödinger values  The Quantum Numbers Explained the arrangement and movement of electrons, spectral lines of poly electronic atoms and gave an acceptable model of an atom. Pauli Exclusion Principle states that no two electrons in the same atom can have identical values for all four of their quantum numbers.  The 4 Quantum Numbers “An Electron’s Address” n Principal Azimuthal Magnetic Spin Quantum Quantum Quantum Quantum Number Number Number Number Specifies the main energy Information about the sub Spin movement of Spatial orientations of an electrons level (orbit) energy level (orbital) orbital Principal Quantum Number (n) Size and Energy of an orbit/shell n=1,2,3,4, … Greater value of n represents bigger orbits with high energies Distance from nucleus also increases Principal Quantum Number Total number of electrons in an orbit = 2n2 Azimuthal Quantum Number Each energy level is divided into sub levels. l defines the shape of sub energy level/orbital. Relationship between n & l l = 0 to (n-1)  EXAMPLE 8 3d 2 n l s→l=0 p→l=1 4f d→l=2 n=4 n=3 f→l=3 l=3 l=2 Magnetic Quantum Number Explains the effect of an orbital in magnetic field i.e. the orientation of an orbital. Orbitals split up in degenerate orbitals (having same energy & size) in a magnetic field. Each degenerate orbital can hold up to 2 electrons Relationship between l & = -l ‣ 0 ‣ +l MAGNETIC QUANTUM NUMBER 3 2p __ __ __ n=2 -1 0 1 l=1 MAGNETIC QUANTUM NUMBER 8 3d __ __ __ __ __ n=3 -2 -1 0 1 2 l=2 MAGNETIC QUANTUM NUMBER 2 4f _ _ _ _ _ _ _ n=4 -3 -2 -1 0 1 2 3 l=3 Aufbau’s Principle states that electrons fill lower-energy atomic orbitals before filling higher-energy ones. Hund’s Rule Every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied and all electrons in singly occupied orbitals have the same spin.  Spin Quantum Number Direction of spin of an electron. Electron which rotates around the nucleus also rotates around its own axis. Either Clockwise (50%) or Anticlockwise (50%) ANTICLOCKWISE CLOCKWISE SPIN QUANTUM NUMBER (ms ) ms = +- 1/2 +1/2 __ __ -1/2  EXAMPLE 3d8 n=3 __ __ __ __ __ l=2 -2 -1 0 1 2 ml = 0 ms = -1/2  LET’S TRY! 4s2 3p5 n= n= l= l= ml = ml = ms = ms =  EXAMPLE 4s2 n=4 __ l=0 0 ml = 0 ms = -1/2  3p5 n=3 l=1 ml=0 -1 0 1 ms= -1/2 Electron Cloud Electron Cloud Electron Cloud Electron Cloud A cloud showing the probability of finding the electron in terms of charged cloud around the nucleus is called Electron Cloud  Atomic Orbitals Atomic orbitals are regions of space where the probability of finding an electron about an atom is highest. s orbital – spherical shapes p orbital – spherical shape d orbital – spherical shapes f orbital – spherical shape  1 STQUARTER PERFORMANCE TASK: Chemical Compound Brochure  1 QUARTER ST PERFORMANCE TASK Make sure to TURN IN your PT in our Google Classroom DEADLINE OF SUBMISSION: ON OR BEFORE SEPTEMBER 20, 2024 (08:00 PM) CHEMICAL BONDING + - Li F  CHEMICAL BONDING It refers to the formation of a chemical bond between two or more atoms, molecules or ions to give rise to a chemical compound These chemical bonds are what keep the atoms together in the resulting compound. VALENCE ELECTRONS AND CHEMICAL BONDS Gilbert Newton Lewis introduced the concept of valence electrons (1946). his idea became the basis for modern theories on chemical bonding. Lewis electron dot structure (LEDS)  VALENCE ELECTRONS electrons which are in the outermost shell or energy level of an atom.  Lewis Electron Dot Symbol (LEDS) Lewis diagrams are graphical representations of elements and their valence electrons. uses dots to represent the number of valence electrons of an element. In a Lewis diagram of an element, the symbol of the element is written in the center and the valence electrons are drawn around it as dots.  Octet Rule atoms tend to form bonds by sharing valence electrons until eight (8) valence electrons surround each atom. Most elements follow the octet rule in chemical bonding, which means that an element should have contact to eight valence electrons in a bond or exactly fill up its valence shell. Having eight electrons total ensures that the atom is stable.  TYPES OF CHEMICAL BONDING  Ionic Bonding bond formed between two ions by the transfer of electrons Formed between a metal and non-metal  Ionic Bonding bond formed between two ions by the transfer of electrons Formed between a metal and non-metal  Elements Metals Ionic bonds form between Metals can be found in the metals and non-metals. middle and on the left hand side of the periodic table. Non-metals Non-metals can be found on the right hand side of the periodic table.  Li F What are ionic bonds? An ionic bond is formed when a metal and non-metal react. Metal atoms become positively charged ions by losing electrons. Non-metal atoms become negatively charged ions by gaining electrons. The oppositely charged ions are very strongly attracted to each other. This is known as an electrostatic attraction.  Forming Forming positive ions negative ions Metal atoms lose electrons to form Non-metal atoms gain electrons to positively charged ions with a full form negatively charged ions with a outer shell of electrons. full outer shell of electrons. + - Li Li F F  Sodium chloride Na Cl Chlorine gains an electron from sodium to become a negative ion (-1). Sodium loses an electron to become a positive ion (+1). Both + - ions now have a full outer shell of electrons and the ionic Na Cl compound sodium chloride is formed.  Li O Li Lithium oxide Each lithium atom loses an electron to become a positively charged ion (1+). The oxygen + 2- atom gains two electrons to become a negatively charged ion 2 Li O (2-).  PROPERTIES OF Ionic Compound + - Li F  Ionic lattice + - + An ionic compound is a regular repeating structure of ions known as - + - a giant ionic lattice. The lattice is composed of a repeating pattern of oppositely charged ions held together + - + by strong electrostatic attractions.  Conduction Ionic compounds can conduct electricity when they are either melted or dissolved in water to form an aqueous solution. In these states, the ions are free to move from place to place. Ionic compounds cannot conduct electricity when solid as their ions are in fixed positions.  Melting and boiling points Ionic compounds consist of oppositely charged ions held together by strong electrostatic attractions. A lot of energy is required to overcome these strong attractions, hence the high melting and boiling points. COVALENT BONDING + - Li F Covalent Bonding F F Covalent Bonding bond formed by the sharing of electrons Covalent Bonding  Elements Covalent bonds form in most non-metal elements and in compounds formed between non-metals. Non-metals Non-metals can be found on the right hand side of the periodic table.  H Cl Molecules O These are examples of covalent molecules. Some are elements H H (substances made of the same type of atom) and some are compounds H H (substances made of two or more types of atom). Cl Cl  Electron Rules Recap He Ne The shells must be filled in order of closest to the nucleus, to furthest from the nucleus. When reacting, the aim is for an atom to achieve a full outer shell. This means the desired electron Two electrons Eight electrons configuration is the same as a noble can occupy the can occupy the gas e.g. like helium and neon shown to first shell. other shells. the left.  H H What are covalent bonds? A covalent bond is formed when two atoms share a pair of electrons. The electrons which contribute towards a covalent bond, are found in the outer shells of the atoms. Usually each atom contributes one electron, but some atoms can react to make multiple covalent bonds.  TWO TYPES OF COVALENT BOND POLAR COVALENT BOND - when electrons are shared but unequally He Ne NON-POLAR COVALENT BOND - when electrons are shared equally  Flourine Each fluorine atom has 7 electrons in the outer shell. Each atom needs to achieve a full outer shell of 8. They F F can each contribute one electron to a covalent bond. Sharing the electrons, means both atoms now have a full outer shell and a simple covalent molecule is made.  Water Water is made of two hydrogen atoms and one oxygen atom. Oxygen has 6 O electrons in its outer shell and needs to achieve 8 to make a full outer shell. Each hydrogen has 1 electron and needs to achieve 2 to have a full shell. Two covalent bonds can be formed to H H make the simple covalent water molecule.  Hydrogen fluoride Fluorine has 7 electrons in its outer shell and needs to achieve 8 to have a F full outer shell. Hydrogen has one H electron. As this electron is in the first shell, hydrogen needs to achieve 2 electrons to have a full shell. The simple covalent molecule of hydrogen fluoride is made by sharing electrons.  PROPERTIES OF MOLECULES F F  Conduction Simple covalent substances can not conduct electricity. This is because charged particles and free movement are required for electrical conduction to occur. Covalent bonds are fixed and the electrons can not move.  Melting Point Covalent bonds are very strong but there are weak intermolecular forces between molecules which do not require a lot of thermal energy to be overcome. This means that simple covalent substance have low melting points and are often liquid or gas at room temperature. Cl Cl  NAMING IONIC AND COVALENT COMPOUNDS CHEMICAL FORMULA Symbolic expression of a compound of a substance. A shorthand of expressing the types and the number of atoms in a substance. CHEMICAL NAME It is a scientific name given to a chemical in accordance with the nomenclature system developed by the INTERNATIONAL UNION OF PURE AND APPLIED CHEMISTRY (IUPAC).  STRUCTURAL FORMULA It is a graphical representation of the molecular structure showing how the atoms are possibly arranged in the real three-dimensional space.  CRISS CROSS METHOD the numerical value of each of the ion charges is crossed over to become the subscript of the other ion. Signs of the charges are dropped.  CRISS CROSS METHOD 1. Write the symbol and charge of the cation (metal) first and the anion (nonmetal) second. 2. Transpose only the number of the positive charge to become the subscript of the anion and the number only of the negative charge to become the subscript of the cation. 3. Reduce to the lowest ratio. 4. Write the final formula. Leave out all subscripts that are 1. IONIC CHARGE  EXAMPLES Write the chemical formula for an ionic compound composed of each pair of ions. 1. the calcium ion and the oxygen ion → CaO 2. the 2+ copper ion and the sulfur ion → CuS 3. the 1+ copper ion and the sulfur ion → Cu2S EXAMPLE: 1. CaO → Calcium oxide 2. NaBr → Sodium bromide 3. MgCl2 → Magnesium chloride Name the following ionic compounds using the Latin System: Auric nitride Silver chloride 3. CuBr Cuprous bromide Name the following ionic compounds using the Latin System: Gold (III) nitride Silver chloride 3. CuBr Copper (I) bromide PRACTICE THIS! 1. KBr →K 1+ , Br 1- → Potassium bromide 2. Na2O →Na 1+ , O 2- → Sodium oxide 3. CaI2 →Ca 2+ , I 1- → Calcium iodide 4. Al4C3 →Al 3+ , C 4- → Aluminum carbide 5. Ca3P2→Ca 2+ , P 3- → Calcium phosphide PRACTICE THIS! 1. FeCl3 →Fe 3+ , Cl 1- → Iron (III) chloride 2. Hg2O →Hg 1+ , O2- → Silver oxide 3. CuSe →Cu 2+ , Se 2- → Copper (II) selenide 4. Ni3N2→ Ni 2+ , N 3- → Nickel (II) nitride 5. CrCl3 → Cr 3+ , Cl 1- → Chromium (III) chloride COVALENT COMPOUNDS EXAMPLE: 1. SO3 → Sulfur trioxide 2. N2O5 → Dinitrogen pentoxide 3. SeBr4 → Selenium tetrabromide Naming Ionic and Covalent compounds