Origin of Elements

Origin of Elements

• Hydrogen and Helium are formed when a dense ball of primordial matter explodes, as per the Big Bang Theory.
• Gravitational forces brought clouds of H and He together, which collapsed into galaxies of billions of stars.
• Fusion reactions form other elements in the centre of stars which requires high temperatures and reactant concentrations.
• Examples of fusion reactions include 24He →8Be , He + Be →12C, and He + C→ 16O.
• Iron (Fe) represents the largest atom resulting from nuclear fusion reactions.
• Elements larger than Iron form from neutrons released by a star's explosion, then captured by Iron nuclei.
• Mainstream fusion reactions produce elements with even mass numbers, while side stream fusion reactions produce odd mass numbers.
• Elements with atomic numbers exceeding 92, known as transuranium elements, undergo laboratory synthesis via fusion reactions.
• An example of laboratory synthesis is Californium-251, the result of carbon-12 nuclei fusing with U-238.

Elemental Composition of Cells

• There are about 30 elements essential for living cells
• Selection factors determining elements in cells include the composition of the Earth, its atmosphere, and the human body
• The elements in the Earth are O2 (48.9%), Fe (18.8%), Si (14.0%), Mg (12.4%), and C (0.10%)
• Sea water and blood plasma feature comparable high concentrations of Cu2+, Mg2+, Ca2+, Na+, and K+

Periodic Table and Cells

• Nearly every periodic table group is represented in living cells
• Cells contain 15 out of the 18 periodic table groups
• Blocks S, P, and D are represented in cells while Block F elements (lanthanides and Actinides) are not

Bulk and Minor Elements

• Major/bulk elements constitute over 99% of cell mass and are present in organic compounds
• Major/bulk elements include Carbon (C), Hydrogen (H), Oxygen (O), Nitrogen (N), Phosphorus (P), and Sulfur (S)
• Minor elements include Mn, Fe, Co, and Ni and are mainly in the fluids bathing cells

Major Elements in Detail

• Major elements are mostly P-block elements
• Major elements can form covalent bonds
• Major elements are non-metals
• Major elements have smaller atomic sizes/numbers
• Major elements are neither overly reactive nor totally inert.

Exclusion of Elements

• Elements with atomic numbers above 92 and 3 below Technetium (Te-43), Astatine (At-85), and Francium (Fr-87) are artificially made
• Rare gases like He, Ne, Ar, Kr, and Xe, and scarce elements like actinides and lanthanides are considered inert
• Lead (Pb2+) induces anaemia by preventing haem synthesis
• Mercury (Hg) causes liver, kidney, and brain damage
• Beryllium (Be2+) interferes with magnesium (Mg2+) functions
• Uranium emits harmful radiations
• Radioactive elements can provide valuable quality of life to cells
• Platinum complex can be an anti-cancer drug

Carbon

• Carbon’s ideal size facilitates stable covalent bond formation
• Carbon has the power of catenation (the ability to bond with itself) which makes straight chains, branches or rings
• Carbon forms multiple bonds with itself or other elements, leading to organic compound diversity
• Carbon is tetravalent

Early Atomic Theories

• Democritus suggested the particulate nature of matter, calling smallest piece "atoms".
• Aristotle proposed matter was continuous and named it "Hyle."
• John Dalton proposed a scientific theory on atom's physical and chemical properties
• Dalton's proposals were based on the law of definite proportion and the conservation of mass

Dalton's Atomic Theory

• Elements consist of indivisible particles known as atoms
• Atoms cannot be created or destroyed
• All atoms of the same element are identical in mass, size, and properties
• Atoms combine in simple ratios to form compounds.
• It is incorrect that atoms are indivisible because of subatomic particles.
• It is incorrect that atoms don't get created because radioactivity and fusion reactions create new atoms.
• It is incorrect that the presence of isotopes are not accounted for because atoms of the same elements can have different masses.
• Hydrogen isotopes include Protium (¹H), Deuterium (²H), and Tritium (³H).

Atomic Models

• J. Dalton: Small, indivisible hard ball model
• J.J. Thompson: Plum pudding or watermelon model
• Ernest Rutherford: Nuclear atom model
• Neil Bohr: Planetary model
• Schrondinger Erwin: Wave mechanic model

Wave Mechanical Model

• Louis de Broglie posited the wave-particle duality of electrons, suggesting they possess both momentum and wavelength
• The uncertainty principle, formulated by Werner Heisenberg, says it is difficult to simultaneously know an electron's position and momentum
• The behavior of electrons in a hydrogen atom is demonstrated through the Schrödinger equation

Atomic Orbitals

• Ψ² (Wave function) is the probability of finding an electron in a small volume.
• Wave functions are referred to as orbitals in atoms.
• An orbital: a volume of space around the nucleus where finding an electron has a 95% probability

Quantum Numbers

• Four quantum numbers include principal, orbital or azimuthal, magnetic, and spin
• The Principle Quantum Number (n) determines the energy level and closeness of electrons to the nucleus.
• Naming for the Principle Quantum starts from K, L, M, etc.
• The maximum number of electrons is 2n², with n representing the energy levels number.
• The Orbital (Azimuthal) Quantum Number (l) indicates the shape of an orbital and ranges from 0 to n-1.
• s-orbital is symmetrical and spherical in shape, and can only contain two electrons, defined as n=1 and l=0.
• The p-orbital is shaped liked a dumb-bell defined as n=2, l=0,1 which includes the s and p orbitals.
• The d shape complex, defined as n=3 and l= 0, 1, 2 which includes the s, p, and d orbitals is more complex
• The Magnetic Quantum Number (mi) refers to the orientation of orbitals in space and includes zero, ranging from -l to +l.
• The s-orbital has no direction of orientation, defined as n=1 l=0 with mi=0.
• For the p-orbital, n=2 with l=0, 1, and mi= -1,0,1
• For the d-orbital, n= 3 with l=0, 1, 2, and mi= -2, -1, 0, +1, +2
• For the f-orbital when n=4 with l=0,1, 2,3, then mi= -3,-2,-1,0,+1,+2,+3
• The Spin Quantum Number (ms) represents the ability of electrons to spin clockwise or anti-clockwise along their axes, ranging from -1/2 to +1/2

Electronic Configuration

• Electronic configuration is the arrangement of electrons in atoms
• Arranging electrons in shells using principal quantum numbers OR using all 4 quantum numbers are the two approaches
• Aufbau's rule is that electron energy levels fill up gradually from the minimum energy level - in all approaches
• Hund's rule is that degenerate atomic orbitals are filled up individually with electrons to the maximum before pairing up, which is done in the second approach
• Pauli's exclusion principle states that that no two electrons in an orbital have all the 4 quantum numbers the same - second approach
• Sodium (Na) follows the octet rule with a configuration of 2, 8, 1 according to the first approach.
• Sodium's configuration (Na) is 1s², 2s², 2p⁶, 3s¹ in the second approach.

Filling Atomic Orbitals

• In the second approach the atomic orbitals in shells fill in the order of 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, etc
• Sc (Scandium) Configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹ 4s²
• Ti (Titanium) Configuration: [Ar] 3d² 4s²
• V (Vanadium) Configuration: [Ar] 3d³ 4s²
• Cr (Chromium) Configuration: [Ar] 3d5 4s1, The 3d and 4s orbitals are half filled this offers special stability
• Mn (Manganese) Configuration: [Ar] 3d5 4s²
• Fe (Iron) Configuration: [Ar] 3d⁶ 4s²
• Co (Cobalt) Configuration: [Ar] 3d⁷ 4s²
• Ni (Nickel) Configuration: [Ar] 3d⁸ 4s²
• Cu (Copper) Configuration: [Ar] 3d¹⁰ 4s¹
• Zn (Zinc) Configuration: [Ar] 3d¹⁰ 4s²

Periodic Table

• Mendeleev created the first periodic table
• Elements are arranged according to atomic numbers, into groups, periods, and blocks
• Groups are vertical arrangements where elements share valence electrons and atomic configurations
• Lithium (Li) atomic configuration is [He] 2s¹
• Sodium (Na) atomic configuration is [Ne] 3s¹
• Potassium (K) atomic configuration is [Ar] 4s¹
• Atomic configurations have one electron outside a noble gas configuration

Groups and Periods

• Most elements have a number of valence electrons and group number that's consistent
• Members of a group with similar physical/chemical properties are "a family"
• Alkali metals (Li, Na, K, Rb, and Cs) are considered family
• Halogens (F, Cl, Br, I) are a family
• Group 4 is not a family like C and Si in this case
• Elements in the same group have varied n-numbers
• Periods are a horizontal arrangement of elements
• Period number is the main energy level that the valence electron occupies
• P1 includes 1s totals 2 elements
• P2 includes 2s, 2p and totals 8 elements
• P3 includes 3s, 3p and totals 8 elements
• P4 includes 4s, 3d, 4p and totals 18 elements
• P5 includes 5s, 4d, 5p and totals 18 elements
• P6 includes 6s, 4f, 5d, 6p and totals 32 elements
• Elements where valence electrons fill the 4f-orbitals have clumsy arrangements, and are in a bottom-row separate row.
• The 4f elements form Lanthanides that start with Cerium, Ce (58) and end at Lutetium, Lu (71).
• Valence electrons fill orbitals which results in placing them below the Lanthanides
• Actinides start at Thorium (Th) 90 and end at Lawrencium (Lr) 103.

Blocks

• Group in which elements occupy valence electrons in same orbital
• G1 and G2 elements that occupy valence electrons within an s-orbital are termed as s-block elements
• Elements starting from Gp3 to Gp0 in the second period contain valence electrons in the p-orbital
• Elements going from group 3 to group 0 are "p-block elements."
• D-block elements, transition elements which involve valence electrons filled d-orbitals
• A 1st transition happens in the 4th period (Sc and Zn)
• A transition happens in the 5th period (Y and Cd)
• A transition happens in the 6th period (L and Hg)
• F block elements with valence electrons filling the f-orbital consist of Lanthanides and Actinides
• S and p blocks are the main group elements

Periodic Properties

• Atomic and physical properties of atoms that reoccur throughout periodic table
• Examples include Atomic/Ionic radius, Ionization energy, Electron affinity, Electronegativity
• Valence electrons are outermost electrons involved in formation of bond
• Core electrons are innermost electrons
• Effective nuclear change is net attraction of proton in the the nucleus for the valence electrons which increases across period decreasing down a column
• Screening effects are decreased attraction of protons for valence electrons by core electrons with trend reversal

Atomic Radius Considerations

• Half the distance between two nuclei bonded in a molecule
• Atom size is dependent on the nucleus attracting valence electrons.
• Screening is constant in a shell but effective nuclear charge raises with an increase in the atomic number
• Smaller atoms results from greater attraction for valence electrons
• Rare gases and halogens increase in atomic radius due to outer shell electron repulsion

Atomic Radius

• Cationic radius is smaller than parent atom
• Lithium cation: Li --- Li⁺
• Sodium cation: Na-----Na⁺
• Isoelectric with noble gas
• Cationic radius decreases over a period while increases down a group
• Anionic radius is larger than parent atom
• Flouride anion: F ----F⁻
• Chloride anion: Cl=====Cl⁻
• Radius of anion (anionic) increases along the period, and reduces down the group.

Ionization Energy

• Ionization energy (Ei): energy to remove loosely bound electrons from gas atoms
• Determinants of Ei: Atomic radius, Shielding effect, Stability of electronic configuration, Penetrative effect
• Down group Ei decreases because atomic radius and screening effect increase, and penetrative effect decreases
• Reverse across period, with similar shielding effect

Exceptions to the trends

• The stability in elctronic configuration offers the best explanation for the following:
• 4Be and 5B differ in that 1s2 2s2 is < than in 1s2 2s2 2p1
• 7N and 80 differ in that 1s2, 2s2, 2p3 is less than in 1s2, 2s2, 2p4
• Applying this idea, each time a 2nd, a 3rd, and more electrons get extracted form one atom, the needed input for ionization of them increases
• Any situation where a noble gas electronic configuration gets involved requires a much more energy required in order to ionize them due to their high stability

Electron Affinity

• Electron affinity (Ea): energy transforms when gas atom has electrons
• Cl (g) + e⁻ ------- Cl⁻(g) : Eea= -348 Kilojoules/mol
• O (g) + e⁻ ------- O⁻ (g): Eea= -141 Kilojoules/mol
• O⁻ (g) + e⁻ ------- O²⁻ (g): Eea= 798 Kilojoules/mol
• Down a column, there is a drop in amount of electron affinity due to high repelling forces for added electrons together.

Trends of Electron Affinity

• The trend is clear with elements more electronegative.
• Increase of electron affinity occurs when reverse crosses throughout periodic table time.
• Lower electron affinities in Fluorine are the discrepancy caused by high repulsion for electrons.
• The determinants effecting electron affinity are atomic radius, configuration stable electronic and anion charge amount.

Electron Affinity in Period 2

• For unexpected values, there may be more electron configuration stability and repulsive electron forces involved.
• B + e⁻→B⁻ = Eea -27KJ, where B is 1s² 2s² 2p¹
• C + e⁻ →C⁻ = Eea -122KJ, where C is 1s² 2s² 2p²
• N + e⁻ →N⁻ = Eea 0KJ, where N is 1s² 2s² 2p³

Electronegativity

• Electronegativity refers to the power of electron of molecules in attracting bond and shared electrons towards itself
• Atoms that attracts many electrons are termed electronegative ones
• The leading factor for electronegativity: radius atom
• More electronegativity is higher in atoms with lower atomic radius
• For electronegativity cases, atoms are measured somehow through relative scale.
• There are no particular units for electronegativity
• Pauling established the reference where 4.0 (Fluorine) were deemed highly electronegative, while 0.7 (Francium) deemed less electronegative
• Elements found to act as very Electronegative are reactive nonmetals (Halogens)
• Elements which act less electronegative are reactive Alkaline metals

Bonding

• Three types of chemical bonds include Ionic/electrovalent, Covalent, and Coordinate Bonds
• Bonds that establish when atoms complete electron transition from one compound to another: Ionic bonds
• Electrostatic pull acts through charged ions forming ionic bond

Ionic Bond Factors & Formation of Bonds

• NaCl (Sodium Chloride) compound is formed in Cycle using Harber and Born processes
• The summation or heat through Hess' theory.
• ∆Hf (heat summation constant in Hess' Law) is, = (∆H sub +Ei)Na + (1/2∆H diss + Eea)Cl + LE or = (∆Hsub + Ei)Na + (1/2 ∆Hdiss - Eea)Cl – LE, depending on compound formed.
• Low ionization positive energy, high electron affinity with negative value, high level- of lattice power results with compound stability
• Negative ∆H more than stable for ionic mixture .
• Formation of ionic mix takes place when there is very difference in electronegativity
• Ionic bonds happen when electronegativity variance is exceeds 1.7
• Covalent polar bonds occur when electronegativity level falls nearly 1.7
• Covalent mixes which stay non-polar tend to happen when very low electronegativity takes place.
• Oxygen molecules in most mixes such as Sodium, Aluminium tend to form more stable ionic and polar bonds.

Examples of Bonding by Electronegativity

• Order of atomic electronegativity Na (1.0), Al (1.5), O (3.5), Cl (2.8), F(4.0), P (2.1), H (2.2)
• Sodium with Oxygen has a of electronegativity Most Ionic character (2.5 variance amount)
• Aluminium, Chlorate will act as polar mix, the covalent bonds (1.3 variance figure).
• Oxygen molecules, fluorine (the nonpole nature by character (0.5 variance figure)
• PH (Phosphorus & Oxide) exhibits a nonpolar nature since lowest (0.1 variance figure .)

Covalent Bonds

• Covalent bond: a way of sharing bonding in atoms as it forms various mix or compounds, H + H → H:H → H₂ = H-H
• Two kinds of Covalent mix bonds exists : Sigma (σ.) & Pi (π)
• Sigma describes in that it's an end to end kind of bonding, it exhibits head-on pull to each mix where their axis ( inter nucleus forms bonding or pull through atomic orbitals
• Pi (π bonding pull comes from sides and the latitudinal to pull them away each others from molecular level as it forms or builds atoms.
• "π ( pi bonds) " more prone to be much fragile and cause reactivity unlike Sigma (σ) bonds as it happens.
• Pi bonding occurs because easily broken and as electrons create available reactions
• Some mixes such as Carbon tetrahydride possess the components from four hydrogen which relates in an equivalent power, a directional state
• More bonds form between ↔Carbon mixtures, which the ethenes are prone by carbon, hydrogen
• A linking which happens when carbon combines among carbon is very unique is a sigma character , with its another pi nature
• Carbon is unequal nature ≠ unlike carbon components of single nature

Hybridization

• Process of mutually blending atomic orbitals with differing shape values and energies, in order to generate equally shaped orbitals, with an appropriate strength and direction
• Methane (CH4), an illustration
• The carbon tends to exhibit divelancy in ground condition. Ground configuration (lowest level)- 2s² ; 2p²
• It exhibits tetravalency, excited state or configuration level 2s¹, 2p³

Sp³ Hybridization

• Process has a nature involving (4) and angle connecting to bond a tetrahedron in an edge (109°eg, CH4-Methane) N/B: In Nitrogen mixes, you will always have 4 components in the atom (central), they surround nature by force
• Has aspects in geometry where Nitrogen is supposed to carry tetrahedron or a Sp³ pull giving in the same kind.
• One set of electrons are left as it'd get repulses by various pairs for other mix which decreases a force. that yields bond for this Ammonia, 107⁰ angular geometry giving the pyramid outlook.

Sp² and SP Hybridization

• SP2 hybridization is a pull where 3 sets of atoms with bond angle of a trigonal look pull the strength approximately 120 g e , the thenes, they all exhibit these three traits also possess an non pulling atomic Sp - where pulling takes at only from atoms in geometry level angle or level (180, that forms ethyne C2H2), they all posses non pulling force as well.

Dative Covalent (Coordinate) Bonding

• This results when bonds from other molecules, atoms where a shared electron is present with donor atom

• Ammonia Hydroxide

• Nitrogen as factor

• 2s2 2p3 where atoms pull one another with equivalent force

• Mix with Hydroxide atom for good: 2s2 2p4 atoms tend to make each other as the same force is used .

• Some more mix bonds exist that builds through coordination but organic features in the bodies

• Mix for Hemoglobin has with it the heme portion in molecule of Ferrous pull bonding

• mix types exist with (Cu(NH3)4)]2+ & [Fe(CN)6)]4

• They each share rule which gives atom transitions and forms shared electrons by force in mix or the geometry to obtain a bonding nature

Exceptions to Octet Rule

• Sodium & Chlorine are Na+ Cl- to make NaCl

• gasses that follow an octet rule

• Most instances the perod with an element and follows carbon , but differs in that it needs atomic orbitals"

• (PCl5 has axial nature to it, where center has in excess to (8 atom).

• Boron Fluoride molecules show to bond only.

        Bonding to form
  

• The configuration for some of them where the atoms exceed the geometry rules :

Lone Pair repulsions

• Examples are Nitrogen mix and Hydroxides
• Molecule in bonding that can't stand and repels has both geometry and other electron pairs
• Lone bond yields big electron intensity that can't withstand a push

Decreasing level of their pull:

• Lone and it's intensity , high intensity pull : Geometry kind geometry that's formed in a geometry by geometry.

VSEPR Theory

• Bond must have electron pair, this leads to minimize the pull in order to make stability in their form.
• Geometry Angles in them:
• -pull from all atoms
• Carbon Mix exhibits at it 120 ° and its geometry
• Mix or is pulled by a geometry in shape.
• The angle is (107°), one of the hydrogen, so this pull makes shape by geometry
• The hydroxides with both has is pulled at (104
• Shape of geometry causes trigonal geometry is shaped by atomic level, with geometry, and hydroxide atomic level, causes angular shape.

Polar Molecules

• The non Polar atom

-has same force, it all repels to attract atom Mixes polar natures . More electronegativity is there is great force to pull every molecule"

• Polar Characteristics :*
• Geometry has has pull it well and cancel geometry
• Molecules pull as
  • (molecules have no Geometry , but do act to be dipole is an intensity

Dipoles exist when there intensity by forces, or geometry in place"

• -OH , water, and are not the kind for zero dipole intensity

• -Mixes tend to produce a higher kind of force all time (homework has (CO2, for mix)