3rd Grade: Matter, Atoms and Structure

• Thema 4 deals with the structure and properties of matter, suitable for 3rd grade students

Atomic Structure

• Matter is composed of atoms, which are extremely small and indivisible particles
• Atoms are minuscule and indivisible
• Atoms are indivisible, massive spheres with constant mass and volume
• Atoms consist of a massive positive nucleus surrounded by moving electrons in a cloud
• Atoms are mostly empty space with a hard, positive nucleus and orbiting electrons
• Atoms are spherical with a central, heavy, positive nucleus where electrons move

Bohr Model

• Electrons orbit the nucleus in specific paths or shells
• Electrons exist in specific energy levels or shells

Electron Configuration

• Electrons have a ground state and excited state
• Heating atoms cause electrons to move further from the nucleus into an excited state
• Excited atoms return to their ground state, releasing excess energy as light (photons) of specific wavelengths

Energy as Wave

• Emitted light's color depends on the energy difference between the excited and ground states of an electron
• A color relates to a specific energy content of radiated light
• Light is an electromagnetic wave described by wavelength (λ), amplitude (A), and speed (c)
• Wavelength is the shortest distance between successive extremes in the wave motion, typically in nanometers (nm)
• Frequency indicates the number of wavelengths passing a point per time unit, such as Hertz (Hz) or s⁻¹
• The order of electromagnetic waves from short to long wavelengths is: Gamma- X rays- ultraviolet- visible light- infrared- microwaves- FM TV radio
• The visible light wavelength ranges from 400 nm (blue) to 700 nm (red)

Emission Spectra

• A line spectrum identifies emitted light from substances when energy is added
• Each element has a unique emission spectrum and serves as a 'fingerprint'
• Only specific distances from the nucleus exist on the specific energy levels

Sommerfeld Model

• Enhanced spectral analysis reveals spectral lines made of closely spaced narrow lines of sharp, principal, diffuse, and fundamental clarity
• Atoms have energy levels divided to sublevels, labelled s, p, d, and f, ranked by energy
• Max electron capacity varies per sublevel: s=2, p=6, d=10, f=14; The # of sublevels per shell matches the shell #, capped at 4

Further Atomic Model Refinement

• Most spectral lines further divide when the emitting gas is placed under a magnetic field
• Additional electron transitions occur under a magnetic field
• Sublevels p, d, and f split to 3, 5, and 7 magnetic levels
• 2 electrons can occupy each magnetic level, with opposite spins

Wave Mechanical Model

• Louis de Broglie states that all moving electrons have wave characteristics (duality principle)
• Werner Heisenberg discovers that determining both the position and speed of an electron is impossible
• Erwin Schrödinger describes the movement of electrons as waves and defines orbitals

Orbitals

• Regions with a 95% probability of finding an electron
• Four types of orbitals are possible: s, p, d, f
• An s-orbital is spherical
• A p-orbital is dumbbell-shaped
• A d-orbital is double dumbbell-shaped or dumbbell-shaped with a ring

Quantum Numbers

• Quantum numbers (n, l, m) give info about the orbital in which the electron exists, quantum number ms gives info about the electron
• Principal quantum # (n) dictates the size of the orbital in space, equals the primary energy level, or shell
• Secondary quantum # (l) :orbital shape, where if l=0: s-orbital (spherical)
• The secondary quantum number is l=1: p-orbital (dumbbell); l=2: d-orbital (various overviews); l =3 : f-orbital (various overviews)
• Magnetic quantum number (m) governs the orientation of the orbital, where m = -l,...,0,...,+l
• Magnetic quantum #: s-orbital 1 orientation - (l = 0, m = 0); p-orbital Three orientations- (l = 1, m = -1, 0, 1)
• Magnetic quantum #: d-orbital , five orientations (l=2 => m = -2,-1,0,1,2); f-orbital = 7 orientations l=3 => m = -3,-2,-1,0,1,2,3
• Spin quantum # (ms) decides the spin direction of the electron: ms = +1/2 (spin up) and ms = -1/2 (spin down)

Pauli Exclusion Principle

• The principle states that no 2 electrons in an atom have the identical values for the 4 quantum numbers, in turn 1 orbital can have 2 electrons
• Two electrons in an orbital are an electron pair, while a single electron in an orbital is an unpaired electron

Electron Configuration

• It details how electrons divide among the differing orbitals
• Orbitals fill with increasing energy so that atoms are stable
• Electron configuration lists the symbols one after another relative to increasing energy
• The exponent is for the number of electrons
• Orbitals overlap, so electrons first fill the 4s-orbital before filling the 3d-orbitals

Spectroscopic Notation

• Full electron configuration notates shell, subshell, and electrons as superscripts
• Shortened electron configuration records the symbol of the previous noble gas, and the orbital
• Block notation (box notation), represents each orbital as a box. Within, arrows represent up or down spins

Hund's Rule

• Electrons in similar orbitals all tend to have the same spin

Stability Rules

• Noble gas configuration - maximum stability
• Filling or half-filling orbitals offer stability
• Electrons can jump from one level to the next to fill the orbitals.

Ion Configuration

• The electronic configurations for ions are the same principles as for neutral atoms
• Cations are electron removal
• Anions are electron addition

Periodic Table and Electron Configuration

• Chemical properties are dependent on electrons on the outermost shell
• Elements in the same column share similar properties
• The PSE is arranged by mass and electron configuration

Organization Of Periodic Table

• The PSE has 7 horizontal rows or periods that relate to the 7 principle energy levels, or shells
• The column corresponds to filling the outer shell
• Elements with the final electron added to a common orbital form a group, which makes those elements share valence electrons and chemical properties

Periodic Groups/Families

• Main group elements are the 'a' groups, and transition metals are 'b' groups
• Elements in the main groups form bonds with either the s-block or p-block
• Number of electrons for similar group labels are the same

s-block Elements

• s-block (last electron to s-sublevel/orbital of the outer shell)
  • group I A: alkali metals (Li to Fr) --> ns¹
  • group II A: alkaline earth metals (Be to Ra) --> ns²

p-block Elements

• p-block (last placed electron in p-sublevel/p-orbital of the outer shell)
  • group III A: earth metals (boron group) --> ns² npx¹ npy`° npz°
  • group IV A: carbon group --> ns² npx¹ npy¹ npz`°
  • group V A: nitrogen group --> ns² npx¹ npy¹ npz¹
  • group VI A: oxygen group --> ns² npx² npy¹ npz¹
  • group VII A: halogens --> ns² npx² npy² npz¹
  • group 0: noble gases --> ns² np⁶

Group 0

• Group 0 possess very stable structure ns² np⁶ aka noble gas setup
• If the element is helium, the arrangement is called a Duet vs Octet

Group B

• The antepenultimate shell (n-1) is supplemented in d-orbitals (d-block), and called transition element

Lanthanides and Actinides

• Lanthanides trail La; and actinides trail Ac, and make-up the f-block, so here the shell (n-2) adds 4f-orbitals to element La, and 5f-orbitals to Ac

Reading The Periodic Table

• It's possible via PSE location, block by shell, to ascertain the configuration

Atomic Properties

• Periodic trends systematically vary in the periodic table qualitatively predicted by the placement of the element

Atomic And Ionic Radii

• Atomic radii grow from top to bottom b/c the # of possessed shells make the atom wider
• The attraction kern-electrons reduces quadratically as atomic radius rises
• Atomic radii descend from L to R because core-electron attraction increases, so radii are smaller nearest right

Ion Size

• +ve charged ions, or cations, bear more narrow radii than corresponding neutral atoms b/c removal makes core attraction better, and radios compress
• -ve charged ions, or anions, sport broader radios than comparable neutral atoms, b/c there become relatively frail core attraction, bigger scatter by electrons, so radio of anion climbs
• -ve is the ion, +grade is the ratio is

Ionization Energy

• Energy req’d to extract the weakest bound electron, relates to core-electron attraction The weaker pull between a core, the smaller the electron can escape, so less fuel must be add’able (smaller energy level 1)

Metals Vs Nonmetals

• Atoms evolve to stable states thru electron emission, or thru catch
• Metals have relatively weak EN and eject
• Nonmetals feature strong EN seize
• The more stable the e’ configuration, the smaller a transfer is; more stable if more small

Oxidation Numbers

• Oxidation numbers parallel a hypothetical charge in a compound where the shift of electrons is complete
• Oxidative states are deduced from the electronic configuration in which the achievement of electronic configuration is the purpose

Explore the structure and properties of matter for 3rd grade, including atoms, the Bohr model, and electron configuration. Learn how atoms consist of a positive nucleus surrounded by moving elections in a cloud. Understand how atoms can be heated where electros move to an excited state.